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Chapter 1Introduction to Chemistry
What is Chemistry?
MatterMatter – anything that has mass and occupies space.
ChemistryChemistry – study of the composition of matter and the changes that matter undergoes.
Because living and nonliving things are made of matter, Because living and nonliving things are made of matter, chemistry affects all aspects of lifechemistry affects all aspects of life
Areas of Chemistry
OrganicOrganic – study of all chemicals containing carbon
See page 8 for examples
InorganicInorganic – study of chemicals that, in general, do not contain carbon. (found mainly in nonliving things)
BiochemistryBiochemistry – study of processes that take place in organisms. (digestion, muscle contraction)
AnalyticalAnalytical – focuses on the composition of matter.
(measuring lead in drinking water)
PhysicalPhysical – area of study that deals with the mechanism, the rate, and the energy transfer that occurs when matter undergoes a change.
Pure & Applied Chemistry
Pure ChemistryPure Chemistry – pursuit of chemical knowledge for its own sake
• Chemists does not expect there to be any immediate practical use for the knowledge
AppliedApplied – research that is directed toward a practical goal or application
In practice, pure chemistry & applied chemistry are linked.
The Scientific Method Logical, systematic approach to the solution of a
scientific problem.
1.1. Making ObservationsMaking Observations – using your senses to obtain information. An observation can lead to a question.
2.2. Making a HypothesisMaking a Hypothesis – a proposed explanation for an observation. A hypothesis is only useful if it accounts for what is actually observed.
3.3. ExperimentExperiment – a procedure that is used to test a hypothesis.
a) Independent variable – a variable that you change during an experiment
The Scientific Methodb) Dependent variable – a variable that is observed
during the experiment.
c) For the results of an experiment to be accepted, the experiment must produce the same result no matter how many times it is repeated or by whom.
4.4. Developing a TheoryDeveloping a Theory – a well-tested explanation for a broad set of observations.
5.5. Scientific LawScientific Law – concise statement that summarized the results of many observations and experiments.
Ex. Gas Laws
Observations Hypothesis Experiments Theory
Scientific Law
The Scientific Method
Summarizes the results of many observations and experiments
A theory is tested by more experiments & modified if necessary
A hypothesis may be revised based on experimental data
Steps do not have to occur in the order shown
Solving Numeric Problems
1.1. Analyze Analyze – identify what is known and what is unknown.
2.2. CalculateCalculate – make the calculations. You may need to convert a measurement or rearrange an equation before you can solve.
3.3. EvaluateEvaluate – after you calculate, evaluate your answer. Is the answer reasonable? Does it make sense?
Chapter 2
Matter
MatterMatter – anything that has mass and takes up space
MassMass – measure of the amount of matter that an object contains
VolumeVolume – measure of the space occupied by the object
Extensive & Intensive Properties
What you observe when you examine a sample of matter is its properties.properties.
1.1. Extensive PropertyExtensive Property – a property that depends on the amount of matteramount of matter in a sample
Ex. Mass, volume, weight, length
2.2. Intensive PropertyIntensive Property – a property that depends on the type of mattertype of matter in a sample (prefix–in means within)
Ex. Hardness, color, odor, luster, conductivity, malleability, ductility, freezing point, boiling point, melting point, density
SubstancesSubstanceSubstance – Matter that has a uniform and definite
composition
• Either an element or a compound
• Also called pure substance
• Rarely found in nature
• Fixed proportions to each other
ExamplesExamples
Diamond Water
Gold Copper
Sugar Nitrogen
MixturesMixtureMixture – a physical blend of two or more
substances that are not chemically combined
• Do not exist in fixed proportions to each other
• Most natural substances are mixtures
• Can usually be separated back into its original components
ExamplesExamples
Concrete Soil
Salt water Milk
Coke Gasoline
Fruit salad Atmosphere
ExamplesExamples
Concrete Soil
Salt water Milk
Coke Gasoline
Fruit salad Atmosphere
Two Types of Mixtures
Homogeneous Mixture (solution)Homogeneous Mixture (solution) – a mixture in which the composition is uniform throughout.
• Consists of a single phase
• Can’t see them separately or separate them physically
ExamplesExamples
stainless steel
air
olive oil
vinegar
Two Types of Mixtures
Heterogeneous Mixture Heterogeneous Mixture – a mixture in which the composition is not uniform throughout.
• Consists of a two or more phases
ExamplesExamples
chicken soup
oil & vinegar mixed
milk
rice crispy treats
Separating Mixtures
Differences in physical properties can be used to separate mixtures
ExamplesExamples
coffee filters
draining pasta
Filtration Filtration – process that separates a solid from a liquid
Separating Mixtures
ExampleExample
separating water from other substances in the water
Distillation Distillation – process of boiling a liquid to produce a vapor and then condensing the vapor into a liquid
States of Matter
1. Solid1. Solid
2. Liquid2. Liquid
3. Gas3. Gas
States of Matter
SolidSolidDefinite shapeDefinite volume
Not easily compressed
CharacteristicsCharacteristics• Does not take the shape of the container• Particles packed tightly together, and often in orderly
arrangement• Almost incompressible• Expands only slightly when heated
States of Matter
LiquidLiquidIndefinite shapeDefinite volume
Not easily compressed
CharacteristicsCharacteristics• Take the shape of the container in which it is placed• Particles in close contact, but arrangement of particles
is not orderly (can flow past each other)• Almost incompressible• Expands slightly when heated
States of Matter
GasGasIndefinite shapeindefinite volume
Easily compressed
CharacteristicsCharacteristics• Take the shape of the container in which it is placed• Can expand to fill any volume• Particles are much farther apart• Easily compressed into a smaller volume
Physical Change
Physical ChangePhysical ChangeSome properties of a material change, but the composition of the material does not change
ExamplesExamples Changes of stateChanges of state such as boiling water,
condensation (boil, freeze, melt, condense)
Physical deformationPhysical deformation such as cutting, denting, stretching, breaking, crushing
Chemical Change
ExamplesExamples Silver spoon tarnishes
Metal rusts Methane burns
Methane burns Sugar ferments
Burn, rot, rust, decompose, ferment, explode, corrode usually mean a chemical change
Chemical ChangeChemical ChangeA change that produces matter with a different composition than the original matter
Elements
Element Element – simplest form of matter that has a unique set of properties.
• cannot cannot be broken down into simpler substances
ExamplesExamples Hydrogen
Nitrogen
Oxygen
Compounds
CompoundCompound – substance that contains two or more elements chemically combined in a fixed proportion.
• Compounds can be broken down into simpler substances by chemical means
ExamplesExamples
Sugar (C12H22O11)
Salt (NaCl)
Water (H2O)
ElementElementSimplest form
CompoundCompound
SubstanceSubstanceDefinite composition
HomogeneousHomogeneousMixtureMixture
Uniform; also calleda solution
HeterogeneousHeterogeneousMixtureMixture
Nonuniform;Distinct phases
MixtureMixtureVariable composition
MatterMatter
Can be separated physically
Can be separatedchemically
Silver Salt StainlessSteel
Cement
Classifying Matter Any sample of matter is either an element, a compound, or a mixture
Symbols Derived From Latin
Sodium Na
Potassium K
Antimony Sb
Copper Cu
Gold Au
Silver Ag
Iron Fe
Lead Pb
Tin Sn
Physical PropertiesPhysical PropertyPhysical Property – a quality or
condition of a substance that can be observed or measuredobserved or measured without changing the substance’s composition
ExamplesExamples
Appearance Density
Texture Malleability
Color Boiling Point
Odor Melting Point
Conductivity Hardness
Chemical Property
ExamplesExamples
Gasoline -- burns in air Iron -- rusts Baking Soda -- reacts with vinegar Copper -- rusts in waterTable salt -- does not react with vinegar
Chemical Property Chemical Property Ability of a substance to undergo a specific chemical change
• Chemical properties can be observed only when a substance undergoes a chemical change.
Recognizing Chemical Changes
ExamplesExamples
Gasoline -- burns in air Iron -- rusts Baking Soda -- reacts with vinegar Copper -- rusts in waterTable salt -- does not react with vinegar
Words such as burn, rot, rust, decompose, ferment, explode, and corrode usually signify a chemical change.
During a chemical change, the compositionof matter always changes.
Recognizing Chemical Changes
Precipitate – Precipitate – solid that forms and settles out of a liquid mixture
Ex.Ex. – ring of soap scum in your bathtub
Possible Clues Possible Clues •Transfer of energy•A change in color•The production of gas•The formation of a precipitate
The only way to be sure a chemical change has occurred is to test the composition of a sample
before and after the change
Law of Conservation of MassDuring any chemical reaction, the mass of the products is
always equal to the mass of the reactants.
ExampleExample
2H2 + O2 2H20
2g 2g = 4g
reactants = product
Chapter 3Chapter 3Observation,Observation,MeasurementMeasurement
and Calculationsand Calculations
Precision and AccuracyPrecision and AccuracyAccuracyAccuracy – measure of how close a measurement comes to the – measure of how close a measurement comes to the actual or actual or truetrue value of whatever is being measured.value of whatever is being measured.
PrecisionPrecision – measure of how close a series of measurements are – measure of how close a series of measurements are to one another.to one another.
Neither accurate nor precise
Precise but not accurate
Precise AND accurate
Determining ErrorDetermining Error
Accepted ValueAccepted Value – – the correct value based on reliable references
Error(can be +or-)=experimental value – accepted value
Percent error = absolute value of error x 100% accepted value
Experimental ValueExperimental Value – – the value measured in the lab
Rules for Counting Rules for Counting SignificantSignificant
FiguresFigures
Nonzero integersNonzero integers always count always count as significant figures. as significant figures.
34563456 hashas
44 sig figs.sig figs.
Rules for Counting Rules for Counting Significant FiguresSignificant Figures
Leading zerosLeading zeros do not count as do not count as significant figuressignificant figures..
0.04860.0486 has has
33 sig figs. sig figs.
Rules for Counting Rules for Counting Significant FiguresSignificant Figures
Zeros at the end of a number Zeros at the end of a number and to the right of a decimal and to the right of a decimal point point are always significant.are always significant.
9.0009.000 has has
44 sig figs sig figs
1.010 1.010 hashas
4 4 sig figssig figs
Rules for Counting Rules for Counting Significant FiguresSignificant Figures
Captive zerosCaptive zeros always count always count as as
significant figures.significant figures.
16.0716.07 has has
44 sig figs. sig figs.
Rules for Counting Rules for Counting Significant FiguresSignificant Figures
Zeros at the rightmost end Zeros at the rightmost end that lie at the left of an that lie at the left of an understood decimal pointunderstood decimal point are are not significant. not significant.
7000 7000 hashas
11 sig fig sig fig
2721027210 has has
44 sig figs sig figs
Rules for Counting Rules for Counting Significant FiguresSignificant Figures
Exact numbersExact numbers have an infinite have an infinite number of significant figures. number of significant figures.
11 inch = inch = 2.542.54 cm, exactlycm, exactly
Rules for Significant Figures in Rules for Significant Figures in Mathematical OperationsMathematical Operations
Multiplication and DivisionMultiplication and Division:: # sig # sig figs in the result equals the number figs in the result equals the number in the in the least precise measurementleast precise measurement used in the calculation. used in the calculation.
6.38 x 2.0 = 6.38 x 2.0 =
12.76 12.76 13 (2 sig figs)13 (2 sig figs)
Rules for Significant Figures Rules for Significant Figures in Mathematical Operationsin Mathematical Operations
Addition and SubtractionAddition and Subtraction: The : The number of decimal places in the number of decimal places in the result equals the number of decimal result equals the number of decimal places in the least precise places in the least precise measurement. measurement.
6.8 + 11.934 = 6.8 + 11.934 =
18.734 18.734 18.7 ( 18.7 (3 sig figs3 sig figs))
Sig Fig Practice #1Sig Fig Practice #1How many significant figures in each of the following?
1.0070 m
5 sig figs
17.10 kg 4 sig figs
100,890 L 5 sig figs
3.29 x 103 s 3 sig figs
0.0054 cm 2 sig figs
3,200,000 2 sig figs
International Systems of UnitsThere are seven SI base units
SI Base Units
Quantity SI base unit SymbolLength Meter m
Mass kilogram kg
Temperature kelvin K
Time second s
Amount mole mol
Luminous intensity candela cd
Electric current ampere A
Metric Prefixes
Meter (m)
Deka(da) 101
Hecto (hm) 102
Kilo (k) 103
Deci (d) 10-1
Centi (c) 10-2
Milli (m) 10-3
Micro (µ) 10-6
Nano (nm) 10-9
Pico (pm) 10-12
Mega (M)
left
right
Other Common Conversions1 cm3 = 1ml
1dm3 = 1L
1 inch = 2.54 cm
1kg = 2.21 lb
454 g = 1 lb
4.18 J = 1 cal
1 mol = 6.02 x 1023 pieces
1 GA = 3.79 L
Units of Lengthmetermeter – the basic SI unit of length or linear measure
Common metric units of length include the centimeter (cm), meter (m), and kilometer (km)
Units of VolumeVolumeVolume -the space occupied by any sample of matter
Volume (cube or rectangle) = length x width x height
The SI unit of volume is the amount of space occupied by a cube that is 1m along each edge. (mm33)
Liter Liter (L) – non SI unit – the volume of a cube that is 10cm along each edge (1000cm1000cm33)
The units milliliter and cubic centimeter are used interchangeably.
1 cm3 = 1ml
1dm3 = 1L
Units of MassCommon metric units of mass include the kilogram, gram,
milligram and microgram.
Weight Weight – is a force that measures the pull on a given mass by gravity.
Weight is a measure of force and is different than mass.
MassMass – measure of the quantity of matter.
Although, the weight of an object can change with its location, its mass remains constant regardless of its location.
Objects can become weightless, but not massless
Units of Temperature
TemperatureTemperature – measure of how hot or cold an object is.
The objects temperature determines the direction of heat transfer.
When two objects at different temperatures are in contact, heat moves from the object at the higher temperature to the object at the lower temperature.
Scientist use two equivalent units of temperature, the degree Celsius and the Kelvin.
Units of TemperatureA change of 1 º on the Celsius scale is equivalent to one
kelvin on the Kelvin scale.
The zero point on the Kelvin scale, 0K, or absolute zero, is equal to -273.15º C.
K = ºC + 273
ºC = K - 273
.
Units of EnergyEnergyEnergy – the capacity to do work or to produce heat.
The joule and the calorie are common units of energy.
The jouleThe joule (J) is the SI unit of energy named after the English physicist James Prescott Joule.
1 calorie1 calorie (cal) - is the quanity of heat that raises the temperature of 1 g of pure water by 1ºC.
1 J = 0.2390 cal
1 cal = 4.184 J
Dimensional AnalysisDimensional analysis – a way to analyze and solve
problems using the units, or dimensions, of the measurements.
How many minutes are there in exactly one week?60 minutes = 1 hour 24 hours = 1 day
7 days = 1 week
1 week 7 days 24 hours 60 minutes = 10,080 min 1 week 1 day 1 hour
1.0080 x 104 min
Dimensional AnalysisHow many seconds are in exactly a 40-hr work week?
60 minutes = 1 hour 24 hours = 1 day7 days = 1 week 60 seconds = 1 minute
40 hr 60 min 60 sec = 144,000 s 1 hr 1 min
1.44000 x 105 s
Dimensional AnalysisGold has a density of 19.3 g/cm3. What is the density in kg/m3
19.3 g 1 kg 1 x 106 cm3 = 1.93 x 104 kg / m3 cm3
1000 g m3
There are 7.0 x 106 red blood cell (RBC) in 1.0 mm3 of blood. How many red blood cells are in 1.0 L of blood?
7.0 x 106 RBC 1 x 106 mm3 1 dm3 = 7.0 x 1012
1.0 mm3 dm3 1 L
DensityIf a piece of led and a feather of the same volume are
weighted, the lead would have a greater mass than the feather.
It would take a much larger volume of feather to equal the mass of a given volume of lead.
Density = mass / volumeD = m / v
Mass is a extensive property (a property that depends on the size of the sample)
Density is an intensive property (depends on the composition of a substance, not on the size of the sample)
QuestionsA student finds a shiny piece of metal that she thinks is
aluminum. In the lab, she determines that the metal has a volume of 245cm3 and a mass of 612g. Was is the density? Is it aluminum?
D = 612g / 245cm3 = 2.50g/cm3
D of aluminum is 2.70 g/cm3; no it is not aluminumA bar of silver has a mass of 68.0 g and a volume of 6.48
cm3. What is the density?
D = 68.0g / 6.48 cm3 = 10.5 g/cm3
Chapter 4Atomic Structure
The AtomYou cannot see the tiny fundamental particles that make up matter.
Yet, all matter is composed of such particles, called atoms
AtomAtom – the smallest particles of an element that retains its identity in a chemical reaction
Several early philosophers and scientists could not observe individual atoms, but still were able to propose ideas on the structure of atoms.
Democritus’s Atomic PhilosophyGreek philospher Democritus (460B.C – 370 B.C.) was
among the first to suggest the existence of atoms.
Democritus believed that matter consisted of tiny, indivisible and indestructible.
• Democritus’s ideas did not explain chemical behavior.
• Lacked experimental support, because his approach was not based on scientific method.
Dalton’s Atomic Theory
According to Dalton’s atomic theory, and element is composed of only one kind of atom, and a compound is composed of particles that are chemical combinations of different kinds of atoms.
1. All elements are composed of tiny indivisible particles called atoms
2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
Dalton’s Atomic Theory3. Atoms of different elements can physically mix together or can
chemically combine in simple whole-number ratios to form compounds.
4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.
Subatomic ParticlesMost of Dalton’s atomic theory is accepted today. Except,
we now know atoms to be divisible.
Atoms can be broken down into smaller particles, called subatomic particlessubatomic particles.
There are 3 kinds of subatomic particles.
1. electrons2. Protons3. neutrons
ElectronsIn 1897, English physicist J.J. Thomson discovered the
electron.
ElectronsElectrons – negatively charged subatomic particles.
Dalton performed experiments that involved passing electric current through gases at low pressure.
Protons and NeutronsAfter a hydrogen atom loses an electron, what is left?
A particle with one unit of positive charge should remain when a typical hydrogen atom loses an electron.
In 1886, Eugene Goldstein observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays.
He concluded they were positive particles. ProtonsProtons – positively charged subatomic particles.
Protons and NeutronsEnglish physicist James Chadwick confirmed the existence
of another subatomic particle.
NeutronNeutron – subatomic particles with no charge but with a mass nearly equal to that of a proton.
Particle Symbol Relative Charge
Relative Mass
Actual mass (g)
electron e- 1- 1/1840 9.11 x 10-28
proton p+ 1+ 1 1.67 x 10-24
neutron n0 0 1 1.67 x 10-24
Rutherford’s Gold-foil ExperimentHowever, the great majority of alpha particles passed
straight through the gold atoms, without deflection.
Also, a small fraction of the alpha particles bounced off the gold foil at very large angles.
Rutherford’s Gold-foil ExperimentBased on his experimental results, Rutherford suggested a
new theory of the atom.
He proposed that the atom is mostly empty space, thus explaining the lack of deflections of most of the alpha particles.
He concluded that all the positive charge and almost all the mass are concentrated in a small region that has enough positive charge to account for the great deflection .
NucleusNucleus – the tiny central core of an atom and is composed of protons and neutrons.
QuestionsDescribe Thomson’s and Millikan’s contributions to
atomic theory.
Thomson – Cathode ray experiments which concluded that electrons must be parts of the atoms of all elements. Millikan determined the charge and mass of the electron.
What experimental evidence led Rutherford to conclude that an atom is mostly empty space?
The great majority of the alpha particles passed straight through the gold foil
Questions
Compare Rutherford’s expected outcome of the gold-foil experiment with the actual outcome.
Expected all alpha particles to pass straight through with little deflection. Found that most passed straight through, but some particles were deflected at large angles and some bounced back.
Distinguishing Among Atoms How are atoms of hydrogen different from atoms of
oxygen?
Elements are different because they contain different number of protons.
Atomic numberAtomic number – of an element is the number of protons in the nucleus of an atom of that element.
Example – all hydrogen atoms have 1 proton and the atomic number of hydrogen is 1.
The atomic number identifies an element.
Distinguishing Among Atoms Most of the mass of an atom is concentrated in its nucleus
and depends on the number of protons and neutrons.
Mass numberMass number – the total number of protons and neutrons in an atom
Example: Helium atom contains 2 protons and two neutrons, so its mass number is 4
If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition.
Distinguishing Among Atoms Example: Oxygen
Atomic number is 8 = number of p+ = e- (So oxygen has 8 electron s and 8 protons.)
Mass number is 16 = number of p+ plus the number of n0. (So oxygen has 8 neutrons)
Number of neutron = mass number – atomic number
197
Au79
Mass number
Atomic number
Isotopes There are some elements that have different kinds of
atoms of the same element
Example – there are three different kinds of Neon atoms
IsotopesIsotopes – are atoms that have the same number of protons, but different numbers of neutrons.
Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.
Isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior.
Chemical Symbols of Isotopes
Write the chemical symbols for three isotopes of oxygen. Oxygen 16, oxygen 17, and oxygen 18.
Mass Number (# protons + # neutrons)
16 17 18
O O O8 8 8
Atomic number (# proton = # electrons)
Atomic Mass The slight difference takes into account the larger masses,
but smaller amounts of the other two isotopes of hydrogen.
Atomic massAtomic mass – of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.
The atomic mass of copper is 63.546 amu. Which of copper’s two isotopes is more abundant: copper -63 or copper-65?
Atomic mass of 63.546 is closer to 63 than 65, thus copper-63 must be more abundant.
Atomic Mass Atomic mass = multiply the mass of each isotope by its
natural abundance, expresses as a decimal, and then add the products.
Element X has two natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with a mass of 11.009 amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.
(10.012 amu x 0.1991) + (11.009 amu x 0.8009) (1.993 amu) + (8.817 amu)
Atomic mass = 10.810
Question
Copper – 63 has a mass of 62.93 amu and 69.2% abundance. Copper-65 has a mass of 64.93 amu and 30.8% abundance. What is copper’s average atomic mass?
(62.93 amu x 0.692) + (64.93 amu x 0.308) (43.548 amu) + (19.998 amu)
Atomic mass = 63.55
Periodic Table
Each element is identified by its symbol place in a square.
The atomic number of the element is shown centered above the symbol. Elements are listed in order of increasing atomic number, from left to right and from top to bottom.
PeriodPeriod - each horizontal row of the periodic table. Within a given period, the properties of the elements vary as you move across it from element to element.
GroupGroup – each vertical column of the periodic table. Elements within a group have similar chemical and physical properties. Each group is identified by a number and the letter A or B.
Chapter 5Models of the Atom
Atomic ModelsRutherford used existing ideas bout the atom and proposed an atomic model in which the electrons move around the nucleus.
However, Rutherford’s atomic model could not explain the chemical properties of element.
Niels Bohr, a student of Rutherford’s, changed Rutherford’s model to include how the energy of an atom changes when it absorbs or emits light.
The Bohr ModelThe Bohr Model – he proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.
The Bohr ModelEach possible electron orbit in Bohr’s model has a fixed energy. The fixed energies an electron can have are called energy levelsenergy levels.
The fixed energy levels of electrons are somewhat like the rungs of the ladder in which the lowest rung of the ladder corresponds to the lowest energy level.
An electron can jump from one energy level to another.
Electrons in an atom cannot be between energy levels.
The Bohr ModelTo move from one energy level to another, an electron must gain or lose jus the right amount of energy.
In general, the higher an electron is on the energy ladder, the farther it is from the nucleus.
A quantum quantum of energy is the amount of energy required to move and electron from one energy level to another energy level.
The energy of an electron is said to be quantized.
The term quantum leap originates from the ideas found in the Bohr model of the atom.
The Quantum Mechanical Model
The Quantum Mechanical ModelQuantum Mechanical Model is the modern description of the electrons in atoms comes from the mathematical solution to the Schrodinger equation.
Like the Bohr model, the quantum mechanical model restricts the energy of electrons to certain values.
Unlike the Bohr model, the quantum mechanical model does not involve an exact path the electron takes around the nucleus.
The quantum mechanical model determines the allowed energies an electron can have an how likely it is to find the electron in various locations around the nucleus
The Quantum Mechanical ModelHow likely it is to find the electron in a particular location is described by probability.
The quantum mechanical model describes of how the electron moving around the nucleus is similar to the motion of a rotating propeller blade.
The propeller blade has the same probability of being anywhere in the blurry regions it produces, but you cannot tells its precise location at any instant.
The Quantum Mechanical Model
The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud.
The cloud is more dense where the probability of finding the electron is high. The cloud is less dense where the probability of finding the electron is low.
It is unclear where the cloud ends, there is at least a slight chance of finding the electron at a considerable distance form the nucleus.
Energy Level
Energy Sublevel ( # = n)
Number of
Orbitals per
Type
Number of
Orbitals per
Level
Number of e- per Sublevel
Max e- in Sublevel
Maximum e- in
Energy Level (2n2)
n = 1 1s 1 1 2e- 2e- 2 e-
n = 22s
2p
1
34
2e-
2e-
2e-
6e-8 e-
n = 3
3s
3p
3d
1
3
5
9
2e-
2e-
2e-
2e-
6e-
10e-
18 e-
n = 4
4s
4p
4d
4f
1
3
5
7
16
2e-
2e-
2e-
2e-
2e-
6e-
10e-
14e-
32 e-
Electron ConfigurationIn most natural phenomena, change proceeds toward the lowest possible energy.
In the atom, electrons and the nucleus interact to make the most stable arrangement possible.
The way in which electrons are arranged into various orbitals around the nuclei of atoms are called electron configuration.electron configuration.
Three rules tell you how to find the electron configurations of atoms.
•The aufbau principle•The Pauli exclusion principle•Hund’s rule
Electron Configuration Rulesaufbau Principle
Electrons occupy the orbitals of lowest energy first.
Pauli Exclusion Principle • An orbital can hold a maximum of 2 electrons. • 2 electrons in the same orbital must have opposite
spins. • An electron is "paired" if it is sharing an orbital with
another electron with an opposite spin. • An electron is "unpaired" if it is alone in an orbital
Paired unpaired
Electron Configuration Rules
Hund’s Rule•Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.•One electron enters each orbital until all the orbitals contain one electron with the same spin direction •For example, three electron would occupy three orbitals of equal energy as follows:•Second electrons then occupy each orbital so that their spins are paired with the first electron in the orbital. Thus each orbital can eventually have two electrons with paired spins.
Electron Configuration PracticeWrite the electron configuration for each atom. How many unpaired electrons does each atom have?
Carbon (atomic number 6 so 6 protons = 6 electrons)
1s22s22p2 2 unpaired electrons
Argon
1s22s22p63s23p6 no unpaired electrons
Silicon
1s22s22p63s23p2 2 unpaired electrons
Exceptional Electron Configurations
Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels.
You can obtain correct electron configurations for the elements up to vanadium (atomic number 23) by following the aufbau diagram for orbital filling.
Cr 1s22s22p63s23p64s23d4 using aufbau
Cr 1s22s22p63s23p64s13d5 correct
Exceptional Electron Configurations
Transition elements are some exceptions to the filling rules.
These exceptions can be explained by the atom’s tendency to keep its energy as low as possible.
These exceptions help explain the unexpected chemical behavior of transition elements.
Shorthand Electron ConfigurationsElectron configurations are often abbreviated by naming the last element with a filled shell (halogens) in brackets and listing only the orbitals after the filled shell.
Na: 1s22s22p63s1
shorthand Na: [Ne] 3s1
Al: 1s22s22p63s23p1
shorthand Al: [Ne] 3s23p1
V: 1s22s22p63s23p6 4s23d3
shorthand V: [Ar] 4s23d3
WavesEach complete wave cycle starts at zero, increases to its highest value, passes through zero to reach its lowest value, and returns to zero again.
AmplitudeAmplitude of a wave is the wave’s height from zero to the crest.
WavelengthWavelength (λ) is the distance between the crests.
Waves
FrequencyFrequency (ν) is the number of wave cycles to pass a given point per unit of time.
The units of frequency are usually cycles per second. The SI unit of cycles per second is called a hertz (Hz)
A hertz can also be expressed as a reciprocal seconds (s-1)
Hz = sHz = s-1-1
Light
The product of frequency and wavelength always equal a constant (c) = the speed of light(c) = the speed of light
c = λν
The wavelength and frequency of light are inversely proportional to each other. As the wavelength increases, the frequency decreases.
According to the wave model, light consists of electromagnetic waves.
Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.
Light
All electromagnetic waves travel in a vacuum at a speed of 2.998 x 102.998 x 1088 m/s m/s
c = 2.998 x 108 m/s
Sunlight consists of light with a continuous range of wavelengths and frequencies. The color of light depends on its frequency.
When sunlight passes through a prism, the different frequencies separate into a spectrumspectrum of color.
A rainbow is an example of this phenomenon.
Each color of the spectrum blends into the next in the order red, orange, yellow green, blue and violet.
In the visible spectrum, red light has the longest wavelength and the lowest frequency.
ElectromagneticSpectrum
Sample ProblemsWhat is the wavelength of radiation with a frequency of 1.50 x 1013 Hz? Does this radiation have a longer or shorter wavelength than red light?
c = λν or λ = c / ν
λ = (2.998 x 108 m/s) / (1.50 x 1013 s-1)λ = 2.00 x 10-5 m (longer wavelength than red light)
What frequency is radiation with a wavelength of 5.00 x 10-8m? In what regions of th e electromagnetic spectrum is this radiation?
c = λν or ν = c / λ
ν = (2.998 x 108 m/s) / (5.00 x 10-8 m)ν = 6.00 x 1015 s-1 (ultraviolet)
When light passes through a prism, the frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrumatomic emission spectrum of the element.
Explanation of Atomic SpectraAtomic line spectra were known before Bohr proposed his model of the H atom. However, Bohr’s model explained why the emission spectrum of H consists of specific frequencies of light.
In the Bohr model, the lone electron in the H atom can have only certain specific energies.
The lowest possible energy of the electron is its ground state. ground state.
In the ground state, the electron’s principal quantum number is 1 (n=1)
Excitation of the electron by absorbing energy raises it from the ground state to an excited state with n = 2,3,4,5…
Explanation of Atomic Spectra
A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level.
The emission occurs in a single abrupt step, called an electronic transition.
Bohr knew from earlier work that the quantum of energy (E) is related to the frequency (ν) of the emitted light by the equation
E = h x E = h x νν
h is the fundamental constant of nature, the “Planck constant” and is equal to 6.626 x 10-34 J·s
Explanation of Atomic Spectra
The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electrons.
Each transition produces a line of a specific frequency in the spectrum.
(Transition to n = 2 energy levelVisible end of the spectra)
(Transition to the n = 1 energy levelUltraviolet part of the spectra)
(Transition to n = 3 energy level, infrared range of spectra)
Quantum MechanicsAlbert Einstein successfully explained experimental data by proposing that light could be described as quanta of energy.
The quanta behave as if they were particles.
Light quanta are called photonsphotons.
Although the wave nature of light was well known, the dual wave-particle behavior of light was difficult for scientists to accept.
Louis de Broglie a French graduate student, asked an important question: Given that light behaves as waves and particles, can particles of matter behave as waves?
The proposal that matter moves in a wavelike way would not be accepted unless experiments confirmed its validity.
Quantum MechanicsGerman physicist Werner Heisenberg examined another feature of quantum mechanics that is absent is classical mechanics.
The Heisenberg uncertainly principleThe Heisenberg uncertainly principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time.
This limitation is critical in dealing with small particles such as electrons.
The Heisenberg uncertainty principle does not matter, however, for ordinary-sized objects such as cars or airplanes.
Recap
The frequency and wavelength of light waves are inversely related. As the wavelength increases, the frequency decreases. (c = λν)
The electromagnetic spectrum consists of radiation over a broad band of wavelengths. The visible light portion is very small. It is in the 10-7 m wavelength rand 1015 Hz (s-1) frequency range.
When atoms absorb energy, electrons move into higher energy levels, and these electrons lose energy by emitting light when they return to lower energy levels.
RecapA prism separates light into the colors it contains. For white light this produces a rainbow of colors. Light from a helium lamp produces discrete lines.
An electron microscope can produce sharp images of a very small object, because of the small wavelength of a moving electron compared with that of light.
The Heisenberg uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time.
Chapter 6The Periodic Table
The Periodic LawMendeleev developed his table before scientists knew about the structure of atoms. He did not know that the atoms of each element contain a unique number of protons.
A British physicist, Henry Moseley, determined an atomic number for each known element.
In the modern periodic table, elements are arranged in order of increasing atomic number.
The Periodic LawThere are seven rows, or periods in the table.
Period 1 has 2 elements, Period 2 has 8 elements, Period 4 has 18 elements & Period 6 has 32 elements.
Each period corresponds to a principal energy level.
There are more elements in higher numbered periods because there are more orbitals in higher energy levels.
The Periodic LawThe elements within a column or group in the periodic table have similar properties.
The properties of the elements within a period change as you move across a period from left to right.
The pattern of properties within a period repeats as you move from one period to the next.
The Periodic LawPeriodic LawPeriodic Law – when elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Group 1 – (alkali metals) are all highly reactive and are rarely found in elemental form in nature
Group 2 – (alkaline earth metals) are silvery colored, soft metals
Group 17- (halogens) the only group which contains elements in all three familiar states of matter at standard temperature and pressure.
Metal, Nonmetals, and MetalloidsThe International Union of Pure and Applied Chemistry (IUPAC) set the standard for labeling groups in the periodic table.
They numbered the groups from left to right 1 – 18,
The elements can be grouped into three broad classes based on their general properties.
• Metals• Nonmetals• Metalloids
Across the period, the properties of elements become less metallic and more nonmetallic.
MetalsAbout 80 % of the elements are metals.
Properties of MetalsProperties of Metals
• Good conductors of heat and electric current.
• Have a high luster or sheen caused by the ability to reflect light
• Solids at room temperature (except Hg)
• Many metals are ductile (can be drawn into wires)
• Most metals are malleable (they can be hammered into thin sheets without breaking)
NonmetalsNonmetals are in the upper-right corner of the periodic table.
There is a greater variation in physical properties among nonmetal than among metals.
Properties of NonmetalsProperties of Nonmetals
• Most are gases at room temperature. S and P are solids, Br is a liquid.
• Nonmetals tend to have properties that are opposite to those of metals.
• In general, nonmetals are poor conductors of heat and electric current. Solid nonmetals tend to be brittle.
MetalloidsThere is a heavy stair-step lines that separates the metals from the nonmetals.
Most of the elements that border this line are metalloids.
Properties of MetalloidsProperties of Metalloids
• Generally has properties that are similar to metals and nonmetals.
• Under some conditions they behave like a metal. Under other conditions they behave like a nonmetal.
QuestionsHow did chemists begin the process of organizing elements?
Used the properties of elements to sort them into groups.
What property did Mendeleev use to organize his periodic table?
In order of increasing atomic mass
How are elements arranged in the modern periodic table?
In order of increasing atomic number
Name the three broad classes of elements.
Metals, nonmetals, and metalloids
Squares in the Periodic TableThe periodic table displays the symbols and names of the elements along with information about the structure of their atoms.
The symbol for the element is located in the center of the square.
The atomic number is above the symbol.
The element name and average atomic mass are below the symbol.
Squares in the Periodic TableThe background colors in the squares are used to distinguish groups of elements. (Ex:2 shades of gold are used for the metals in Groups IA and 2A)
Group IA elements are called alkali metalsalkali metals. Group 2A elements are called alkaline earth metalsalkaline earth metals.
The nonmetals of Group 7A are called halogens.halogens.
Group 8A elements are called Noble GasesNoble Gases
Groups 1B – 8B are called transition metals transition metals
The two periods usually located at the bottom of the periodic table separate from the main table are called inner inner transition elements.transition elements. Period 8 is called the Lanthanide Lanthanide SeriesSeries and Period 9 is called the Actinide SeriesActinide Series
Electron Configuration in Groups
Electrons play a key role in determining the properties of elements.
So there is a connection between an element’s electron configuration and its location in the periodic table.
Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals based on their electron configurations.
The Noble Gases are in Group 8A and are sometimes called inert gases because they rarely take part in a reaction.
Electron Configuration in Groups
The highest occupied energy level for each element, (the s & p sublevels) are completely filled with electrons.
Helium (He) 1s2
Neon (Ne) 1s22s22p6
Argon (Ar) 1s22s22p63s23p6
Krypton (Kr) 1s22s22p63s23p63d104s4s224p4p66
s sublevel p sublevel
The Representative ElementsElements in groups 1A through 7A are often referred to as representative elementsrepresentative elements because they display a wide range of physical and chemical properties.
In atoms of representative elements, the s and p sublevels of the highest occupied energy level are not filled.
Lithium(L) 1s22s1
Sodium (Na) 1s22s22p63s1
Potassium (K) 1s22s22p63s23p64s4s11
s sublevel
The Representative Elements
In atoms of carbon, silicon, and germanium, in Group 4A, there are four electrons in the highest occupied energy level
For any representative elements, its group number equals the number of electrons in the highest occupied energy level.
Carbon (C) 1s22s22p2
Silicon (Si) 1s22s22p63s23p2
Germanium (Ge) 1s22s22p63s23p64s4s223d104p4p22
p sublevels sublevel
Transition MetalsElements in the B groups are referred to as transition transition elements. elements.
There are two types of transitions elements: transition transition metalsmetals and inner transition metalsinner transition metals
In atoms of a transition metal, the highest occupied s sublevel and a nearby d sublevel contain electrons.
These elements are characterized by the presence of electrons in d orbitals.
IonsSome compounds are composed of particles called ions. An ionion is an atoms or group of atoms that has a positive or negative charge.
An atom is electrically neutral because it has equal numbers of protons and electrons.
Positive and negative ions from when electrons are transferred between atoms.
Atoms of metallic elements tend to form ions by losing one or more electrons from their highest occupied energy levels.
A sodium atom tend to lose one electron.
CationsIn the sodium ion, the number of electrons (10) is no longer equal to the number of protons (11).
Because there is more positively charged protons than negatively charged electrons, the sodium ion has a net positive charge.
An ion with a positive charge is called a cationcation.
The charge for a cation is written as a number followed by a plus sign. (Example: 1+ )
If the charge is 1+, the number 1 is usually omitted from the complete symbol for the ions. (Na+)
AnionsAtoms of nonmetallic elements, such as chlorine, tend to form ions by gaining one or more electrons.
A chlorine atom tend to gain one electron.
In a chlorine ion, the number of electrons (18) is no longer equal to the number of protons (17).
Because there are more negatively charged electrons than positively charged protons, the chloride ion has a net negative charge.
An ion with a negative charge is called an anionanion.
Examples: Cl-, S2-
Trends in Ionization Energy
Recall that electrons can move to higher energy levels when atoms absorb energy.
Sometimes there is enough energy to overcome the attraction of the protons in the nucleus.
The energy required to remove an electron from an atom is called ionization energyionization energy.
The energy to remove the first electron from an atom is called the first ionization energyfirst ionization energy.
The cation produced has a 1+ charge.
Ionization Energy
The energy to remove the first electron from an atom is called the first ionization energyfirst ionization energy. The cation produced has a 1+ charge.
The second ionization energysecond ionization energy is the energy required to remove an electron from an ion with a 1+ charge. The ion produced has a 2+ charge.
The third ionization energythird ionization energy is the energy required to remove an electron from an ion with a 2+ charge. The ion produced has a 3+ charge.
Trends in ElectronegativityThere is a property that can be used to predict the type of bond that will form during a reaction.
This property is electronegativitylectronegativity, which is the ability of an atom of an element to attract electrons when the atom is in a compound.
In general, electronegativity values decrease from top to bottom within a group.
For representative elements, the values tend to increase from left to right across a period.
Electronegativity increases
Nuclear charge increases
Shielding is constant
Ionization energy increases
Atomic size decreases
Atomic size increases
Ionic size increases
Ionization Energy decreases
Electronegativity decreases
Nuclear charge increases
Shielding increases
Size of cation decreases
Size of anions decreases
Trends for Groups 1A Through 8A• Can be explained by variations
in atomic structure• Increase in nuclear charge
within groups & across periods, also shielding within groups
Chapter 7Ionic and Metallic Bonding
valence ElectronsScientists learned that all of the elements within each group of the periodic table behave similarly because they have the same number of valence electrons.
valence electronsvalence electrons are the electrons in the highest occupied energy level of an element’s atom.
The number of valence electrons largely determines the chemical properties of an element.
To find the number of valence electrons in an atom of a representative elements, simply look at its group number
Elements of Group IA have one valence electron. Elements in Group 4A have four valence electrons, and so forth
valence ElectronsThe noble gases, Group 8A, are the only exceptions to the group-number rule.
Helium has two valence electrons, and all of the other noble gases have eight.
valence electrons are usually the only electrons used in chemical bonds.
As a general rule, only the valence electrons are shown in electron dot structures.
Electron dot structures are diagrams that show valence electrons as dots.
Electron Dot Structures
The Octet RuleNoble gases, such as neon and argon, are unreactive in chemical reactions. (They are stable)
Gilbert Lewis explained why atoms form certain kinds of ions and molecules in the octet rule
The Octet RuleThe Octet Rule - in forming compounds, atoms tend to achieve the electron configuration of a noble gas. An octet is a set of eight. (each noble gas except helium has eight electrons in its highest energy level)
Atoms of the metallic elements tend to lose their valence electrons, leaving a complete octet in the next-lowest energy level. Atoms of some nonmetallic elements tend to gain electron or to share electrons with another nonmetallic element to achieve a complete octet.
Formation of CationsUsing electron dot structures, you can show the ionization of some elements more simply.
Na· Na+ + e- Sodium atom Sodium ion electron neutral 1 unit of + charge 1 unit of - charge
·Mg· Mg2+ + 2e- Magnesium atom Magnesium ion electron neutral 2 unit of + charge 2 units of - charge
Transition MetalsFor transition metals, the charges of cations may vary.
An atom of iron (Fe) may lose two, or three electrons forming either Fe2+ or Fe3+ ions.
Some ions formed by transition metals do not have noble gas electron configurations and are therefore exceptions to the octet rule.
Ag is an example - 1s22s22p63s23p63d104s4s224p4p664d105s5s11
To achieve the structure of krypton, which is the preceding noble gas, a silver atom would have to lose eleven electrons.
Transition MetalsIons with charges of three or greater are uncommon, and losing eleven electrons is highly unlikely.
If Ag loses its 5s1 electron, the configuration that results, (4s4s224p4p664d10) with 18 electrons in the outer energy level and all of the orbitals filled, is relatively favorable in compounds.
Such a configuration is known as pseudo noble-gas electron pseudo noble-gas electron configuration. configuration.
Ag forms a positive ion (Ag+) in this way.
Formation of AnionsThe gain of negatively charge electrons by a neutral atom produces an anion.
The name of an anion of a nonmetallic element is not the same as the element name. The name of the ion typically ends in -ide.-ide.
Chlorine atom (Cl) forms a chloride ion (Cl-)
Oxygen atom (O) forms an oxide ion (O2-)
Because they have relatively full valence shells, atoms of nonmetallic elements attain noble-gas electron configurations more easily by gaining electrons than by losing them.
Formation of AnionsChlorine belongs to Group 7A and has seven valence electrons. A gain of one electron gives chlorine an octet and converts a chlorine atom into a chloride ion.
Atoms of nonmetallic elements form anions by gaining enough valence electrons so as to attain the electron configuration of the nearest noble gas.
The chloride ion has the same electron configuration as the noble gas argon.
Chloride ion (Cl-) 1s22s22p63s23p6
Argon (Ar) 1s22s22p63s23p6
Formation of Ionic CompoundsCompounds composed of cations and anions are called ionic ionic compoundscompounds.
Ionic compounds are usually composed of metal cations and nonmetal anions. Ex: NaCl is formed from Na+ + Cl-
Although they are composed of ions, ionic compounds are electrically neutral. The total + charge of the cations equals the total – charge of the anions.
Anions and cations have opposite charges and attract one another by means of electrostatic forces.
The electrostatic forces that hold ions together in ionic compounds are called ionic bondsionic bonds.
Formation of Ionic CompoundsLook at the reaction of a Na atom and a chlorine atom.
Na has 1 valence electron that it can easily lose. (Na is in group 11A of the representative elements, thus has 1 1 valence electron)
Cl has seven valence electrons and can easily gain one electron. (Cl is in group 77A of the representative elements, thus has 77 valence electrons)
If Na loses its valence electron it achieves the stable electron configuration of neon. If Cl gains a valence electron, it achieves the stable electron configuration of argon. (Remember the Octet Rule)
Formation of Ionic CompoundsWhen Na and Cl react, the Na atom gives its one valence electron to a Cl atom. They react in a 1:1 ratio and both ions have stable octets.
+
Na+ Cl-
1s22s22p6 1s22s22p63s23p6
Formula UnitsChemists represent the composition of substances by writing
chemical formulas. A chemical formulachemical formula shows the kinds and numbers of atoms in the smallest representative unit of a substance.
NaCl is the chemical formula for sodium chloride.
A Formula unitFormula unit is the lowest whole-number ratio of ions in an ionic compound. One Na+ to each Cl-, thus the formula unit for sodium chloride is NaCl.
Even though ionic charges are used to derive the correct formulas, they are not shown when you write the formula unit of the compound
Formula UnitsThe ionic compound Magnesium chloride (MgCl2) contains
magnesium cations (Mg2+) and chloride anions (Cl-)
In MgCl2, the ratios of Mg2+ to Cl- is 1:2 (One Mg2+ to two Cl-). Its formula unit is MgCl2
Because there are twice as many Cl- (each with a 1- charge) as Mg2+ (each with a 2+ charge), the compound is electrically neutral.
Another example: Al3+ + Br- combine to form AlBr3.
Metallic Bonds & PropertiesMetals are made up of closely packed cations rather than neutral
atoms.
The valence electrons of metal atoms can be modeled as a sea of electrons. (they are mobile and can drift freely from one part of the metal to another).
Metallic bondsMetallic bonds consists of the attraction of the free-floating valence electrons from the positively charged metal ion.
The sea-of-electrons model explains many physical properties of metals. – Good conductors of electrical current because electrons
can flow freely.– Ductile – they can be drawn into wires.– Malleable – they can be hammered or forced into shapes.
Crystalline Structure of MetalsThere are several closely packed arrangements that are possible.
• body-centered cubic arrangement• face-centered cubic arrangement• hexagonal close-packed arrangement
Body-centered cubicBody-centered cubicEvery atom (except those on theSurface) has eight neighbors.
Crystalline Structure of MetalsFace-centered Face-centered cubic arrangement
• every atom has twelve neighbors.
Crystalline Structure of MetalsHexagonal close-packedHexagonal close-packed arrangement
• every atom also have twelve neighbors. Because of the hexagonal shape, the pattern is different from the face-centered.
Alloys
Very few of the metallic items that you use every day are pure metals. Ex: spoons.
Most of the metals you encounter are alloys.
Alloys are mixtures composed of two or more elements., at least on of which is a metal. Ex: Brass (Cu & Zn)
Alloys properties are often superior to those of their component elements.
Sterling silver (92.5% silver & 7.5% copper) is harder and more durable than pure silver, but still soft enough to be made into jewelry and tableware.
Laws Governing Formulas & NamesLaw of Definite ProportionsLaw of Definite Proportions
A chemical formula tells you (by subscripts) the ratio of atoms of each element in the compound.
Ratios of atoms can also be expressed as ratios of masses.
100 g of MgS breaks down into 43.12g Mg and 56.88g of sulfur.
100g MgS 1 mol MgS 1 mol Mg 24.305g Mg = 43.12g Mg
56.4g MgS 1 mol MgS 1 mol Mg
100g MgS 1 mol MgS 1 mol S 32.06g S = 56.88g S 56.4g MgS 1 mol MgS 1 mol S
Chapter 10Chemical Quantities
Measuring MatterAvogadro’s number is the number of representative particles in a mole, 6.02 x 1023.
The term representative particles refers to the species present in a substance: usually atoms, molecules or formula units.
Representative particles for ionic compounds is the formula unit : CaCl2 , NaCl
Representative particles for molecular compounds is the molecule: H2O , H2
Representative particles for most elements is the atom: Fe, Li
Measuring MatterA mole of any substance contains Avogadro’s number of representative particles or 6.02 x 1023 representative particles.
The relationship, 1 mole = 6.02 x 1023 representative particles, is the basis for a conversion factor to convert numbers of representative particles to moles.
How many moles of Mg is 1.25 x 1023 atoms of Mg?
1.25 x 1023 atoms Mg (1 mol Mg / 6.02 x 1023 atoms Mg)
Measuring MatterHow many atoms are in 2.12 mol of propane (C6H8)?
In the formula of a molecule of C3H8 , the subscripts show that propane is composed of 14 atoms: 3 atoms of C and 8 atoms of H.
2.12 mol C6H8 6.02 x 1023 molecules C6H8 11 atoms 1 mol C6H8 1 molecule of C6H8
1.40 x 1025 atoms
Mass of a MoleThe atomic mass of an element (mass of a single atom) is expressed in atomic mass units (amu)
The atomic masses are relative values based on the mass of the most common isotope of carbon 12.
The atomic mass of an element expressed in grams is the mass of a mole of the element.
The mass of a mole of an element is its molar mass.
Molar mass of C is 12.0 g. H – 1.0 g, S – 32.1g
Molar mass is the atomic mass of an element rounded off to the first decimal place.
Molar MassIf you were to compare 12.0g of C atoms with 16.0g of O atoms, you would find they contain the same number of atoms.
The molar mass of any element contains 1 mole or 6.02 x 1023 atoms of that element.
12.0g of C is 1 mol of C atoms
1.0 g of H is 1 mol of H atoms
Molar mass is the mass of 1 mole of atoms of any element.
Mass of a Mole of a CompoundTo find the mass of a mole of a compound, you must know the formula of the compound.
A molecule sulfur trioxide, SO3, is composed of one atom of sulfur and three atoms of oxygen.
Calculate the mass of a molecule of SO3 by adding the atomic masses of the atoms making up the molecule.
The atomic mass of Sulfur is 32.1g and the mass of three Oxygen atoms is 48.0g (3 x 16.0), so the molecular mass of SO3 is 80.1g (32.1 + 48.0)
The molar mass of any compound is the mass of 1 mole of that compound.
Mass of a Mole of a Compound
1 mole of SO3 has a mass of 80.1g and is the mass of 6.02 x 1023 molecules of SO3
To calculate the molar mass of a compound, find the number of grams of each element in one mole of the compound and then add the masses of the elements.
The method for calculating molar mass applies to any compound, molecular or ionic.
Mole/Mass RelationshipYou need 3.00 mol of NaCl. How do you measure this amount? What mass in grams is 3.00 mol of NaCl?
3.00 mol NaCl 58.5 g NaCl = 176g NaCl (use the molar mass)
1 mol NaCl
When you measure 176g of NaCl on a balance, you are measuring 3.00 mol of NaCl.
What is the mass of 9.45 mol of aluminum oxide? (Al2O3)
9.45 mol Al2O3 102.0g Al2O3 = 964 g Al2O3
1 mol Al2O3
Mole/Mass RelationshipHow many moles of sodium sulfate (Na2SO4) is in 10 g of Na2SO4?
10.0 g Na2SO4 1 mol Na2SO4 = 7.04 x 10-2 mol Na2SO4
142.1 g Na2SO4
How many moles of iron(III) oxide are contained in 92.2 g of pure Fe2O3?
92.2 g Fe2O3 1 mol Fe2O3 = 0.578 mol Fe2O3
159.6 g Fe2O3
Mole/Volume RelationshipThe volume of one mole of different solid and liquid substances are not the same. However, the volumes of moles of gases measured under standard condition are much more predictable.
Avogadro’s hypothesis states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles.
If you buy a party balloon filled with helium and take it home on a cold day, you might notice that the balloon shrinks while it is outside.
The volume of a gas varies with a change in temperature.
Mole/Volume RelationshipThe volume of a gas also varies with a change in pressure. An increase in pressure causes the volume of the gas to decrease.
Because of these variation due to temperature and pressure, the volume of a gas is usually measured at standard temperature and pressure.
Standard temperature and pressure (STP) means a temperature of 0ºC and a pressure of 101.3 kPa (1atm)
At STP, 1 mole or 6.02 x 1023 representative particles of any gas occupies 22.4L.
22.4 L is called the molar volume of gas.
Mole/Volume RelationshipIf you have 0.375 mol of O2 gas, what volume at STP will this gas occupy?
0.375 mol O2 22.4L O2 = 8.40 L O2
1 mol O2
Determine the volume in liters of 0.60 mole of SO2 gas at STP.
0.60 mol SO2 22.4L SO2 = 13 L SO2
1 mol SO2
How many moles of H2 are in 0.200 L at STP?
0.200 L H2 1 mol H2 = 8.93 x 10-3 mol H2
22.4 L H2
Molar Mass From DensityDifferent gases have different densities.
Usually the density of a gas is measured in grams per liter (g/L)
The density of a gas at STP and the molar volume at STP can be used to calculate the molar mass of the gas.
The density of a gaseous compound containing C and O is 1.964 g/L at STP. What is the molar mass of the compound?
1.964 g 22.4 L = 44.0 g/mol L 1 mol
Percent CompositionThe relative amounts of the elements in a compound are expressed as the percent composition or the percent by mass of each element in the compound.
The percent composition of a compound consists of a percent value for each different element in the compound.
The percent composition of K2CrO4 is K = 40.3%, Cr = 26.8%, O = 32.9%. (They must total 100%)
The percent by mass of an element in a compound is the number of grams of the element divided by the mass in grams of the compound, multiplied by 100%.
Percent Composition% mass of element = mass of element x 100% mass of compound
When a 13.60 g sample of a compound containing only Mg and O is decomposed, 5.40g of O is obtained. What is the percent composition of this compound?
% O = 5.40 g / 13.60g x 100% = 39.7%
% Mg = 13.60 g – 5.40 g / 13.60g x 100% = 60.3%
Percent Composition by Formula% mass = mass of element in 1 mol compound x 100% molar mass of compound
Calculate the percent composition of propane C3H8
% C = 36.0 g / 44.0 g x 100% = 81.8%
% H = 8.0 g / 44.0 g x 100% = 18.0%
Percent Composition as a Conversion Factor
How much C and H are contained in 82.0 g of propane? (C3H8)
Calculate the percent composition of propane C3H8
% C = 36.0 g / 44.0 g x 100% = 81.8%
% H = 8.0 g / 44.0 g x 100% = 18.0%
In a 100 g sample of propane you would have 81.8 g of C and 18 g of O.
(82.0 g propane)(81.8 g C / 100 g propane) = 67.1 g C
(82.0 g propane)(18 g O / 100 g propane) = 15 g H
