CHEMISTRY - BROWN 13E CH.2 - ATOMS, MOLECULES & IONSlightcat-files.s3.amazonaws.com/packets/admin... · Atoms Moles Grams EXAMPLE: Determine the mass (in grams) found in 7.28 x 1028

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CHEMISTRY - BROWN 13E

CH.2 - ATOMS, MOLECULES & IONS

CONCEPT: GROUP NAMES AND CLASSIFICATIONS

Ever wonder where did this periodic table ever come from?

• At the end of the 18th century, Lavoisier compiled a list of the 23 elements known at the time.

• In 1869, Dmitri Mendeleev coined the term “Periodic Table” .

• Today the total is 114 and still counting!

Now, to understand chemistry fully it will be imperative that you memorize and learn the different portions of the Periodic

Table.

Phase Differences

At room temperature (between 20 oC to 25oC), all elements are _______________ except:

• Mercury and bromine are _______________ . • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine and the Noble Gases are ____________.



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CONCEPT: CHARGE DISTRIBUTIONS OF THE PERIODIC TABLE

A majority of the elements on the periodic table are reactive because they all want to be like the _____________________.

• They have the perfect number of electrons in their outer atomic shells.

1. Metals tend to __________ electrons to become positively charged ions called _____________.

• Metals that have ONLY one charge are referred to as ________________ metals.

• Metals that have MORE THAN one charge are referred to as ________________ metals.

2. Nonmetals tend to __________ electrons to become negatively charged ions called _____________.



Page 3

CONCEPT: ELEMENT SYMBOLS

Some of the names and symbols for the elements are easy to recognize like Aluminum is Al, but some others aren’t.

EXAMPLE 1: Identify the elements by their given symbols.

a. Au b. Hg c. Pb d. Fe e. Ag

Some elements exist in nature connected to their exact double.

We call these chemical Siamese twins ________________________________.

To recall them just remember this funny phrase:

Have No Fear Of Ice Cold Beer

Some elements exist in nature as monoatomic elements such as _______________ & _______________.

Some elements exist in nature as polyatomic molecules such as _______________ & _______________.



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CONCEPT: MASS CONVERSIONS

The _______________ is the chemical unit for the amount of a substance.

One mole (1 mol) contains 6.022 x 1023 entities, which is known as _________________________________________.

• Entities means ______________________ , ______________________ or ______________________.

• We use ______________________ when dealing with a single, individual element.

• We use _________________ or _________________when dealing with more than one element or a compound.

6.022 x 1023 atoms of Fe is equal to 1 mole of Fe and has a mass of 55.85 amu

Atoms Moles Grams

EXAMPLE: Determine the mass (in grams) found in 7.28 x 1028 nitrogen atoms.

6.022 x 1023 molecules of H2O is equal to 1 mole of H2O and has a mass of 18.016 amu

Molecules Moles Grams

EXAMPLE: Determine how many molecules of carbon dioxide, CO2, are found in 75.0 g CO2.



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CONCEPT: MASS CONVERSIONS (PRACTICE)

PRACTICE 1: If the density of water is 1.00 g/mL at 25oC calculate the number of water molecules found in 1.50 x 103 µL

of water.

PRACTICE 2: Calculate the number of oxygen atoms found in 783.9 g CuSO4 · 5 H2O.

PRACTICE 3: The density of the sun is 1.41 g/cm3 and its volume is 1.41 x 1027 m3. How many hydrogen molecules are in

the sun if we assume all the mass is hydrogen gas?

PRACTICE 4 (CHALLENGE): A cylindrical copper wire is used for the fences of a house. The copper wire has a diameter

of 0.0750 in. How many copper atoms are found in 5.160 cm piece? The density of copper is 8.96 g/cm3. ( V = π · r2 · h ).



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CONCEPT: ATOMIC MASS

Whether you call it atomic mass or weight both terms tell us the combined mass of the protons and neutrons in an element.

• The atomic masses listed for the elements on the periodic table are the ____________________ of their isotopes.

• Isotopes are elements with the _______________ number of protons, but _______________ number of neutrons.

Atomic Mass = [(Mass of Isotope 1) x (Fractional Abundance 1)] + [(Mass of Isotope 2) x (Fractional Abundance 2)]

EXAMPLE 1: Antimony has two common isotopes. If one of the isotopes 121Sb has an isotopic mass of 120.9038 amu and

a natural abundance of 57.25%, what is the isotopic mass (to 4 significant figures) of the other isotope? The atomic mass of

antimony is 121.8 g/mol.

EXAMPLE 2: The atomic mass of an imaginary element A is 251.7 amu. If element A consists of two isotopes that have

atomic masses of 250 and 253 respectively, what is the natural abundance of each isotope?



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CONCEPT: MASS SPECTROMETRY

Mass spectrometry involves the __________________, __________________, and __________________ of gaseous ions according to their mass to charge ratios.



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CONCEPT: STRUCTURE OF THE ATOM

We learned that the basic functional unit in chemistry is the _____________ .

• Now it’s time to go into an atom to figure out its components: _________ subatomic particles.

In the center of an atom there is the _____________,

• It contains the subatomic particles: _____________ and _____________.

• Spinning around it we find the third subatomic particle: the _____________.

• PROTONS are _________________ charged subatomic particles.

• ELECTRONS are _________________ charged subatomic particles.

• NEUTRONS are _________________ charged subatomic particles.

ATOMIC NUMBER equals the number of _________________ and determines _________________ of an element.

ATOMIC MASS equals the number of ____________________________ in an element.

EXAMPLE: Identify the unknown element.

a. Element X (8 protons, 8 electrons, 8 neutrons) b. Element Y (35 protons, 36 electrons, 46 neutrons) c. Element Z (12 protons, 10 electrons, 13 neutrons)

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CONCEPT: MODERN ATOMIC THEORY

According to the Law of ________________________________ in a reaction matter is neither created nor destroyed.

• Originated in 1789 by Antoine Lavoisier.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

According to the Law of ________________________________ all samples of a compound, no matter on their origin or

preparation has the same ratio in terms of their elements.

• Originated in 1797 by Joseph Proust.

Mass Ratio = 12.0gC)((32.0gO)

= 0.375

According to the Law of ________________________________ when two elements (A & B) form different compounds, the

masses of element B that combine with 1 g of A are a ratio of whole numbers.

• Originated in 1804 by John Dalton.

Mass Ratio = 16.0gO)((14.0gN )

=1.143

Mass Ratio = 32.0gO)((14.0gN )

= 2.286

The ratio of the two mass ratios obtained then gives us a whole number:

2.2861.143

= 2.0

CO2

NO

NO2



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CONCEPT: MODERN ATOMIC THEORY (PRACTICE)

EXAMPLE 1: A 15.39 g sample of iodine reacts with 62.92 g of chlorine to form iodine pentachloride, ICl5. If iodine

pentachloride is the only product formed calculate its mass.

EXAMPLE 2: Two samples sodium fluoride decompose into their constituent elements. The first sample produces 15.8 kg

of sodium and 20.1 kg of fluorine. If the second sample produces 192.0 g of sodium, how many grams of fluorine were also

produced?

PRACTICE: Which of the following is an example of the law of multiple proportions?

a. A sample of bromine (Br) contains equal amounts of its two isotopes.

b. Two different samples of H2O have the same mass ratio.

c. The atomic mass of sodium (Na) is 22.99 amu.

d. Two different compounds composed of sulfur (S) and oxygen (O) have different mass ratios: 2.48 g O: 1 g S and

1.24 g O: to 1 g S.



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CONCEPT: THOMSON CATHODE RAY TUBE EXPERIMENT

J.J. Thomson’s cathode ray tube experiments led to the discovery of the _____________________ .

Apply an Electric Field When an electric field is applied across the cathode ray tube, the cathode ray is attracted to the plate with a ______ charge.

Applying a Magnetic Field A moving charged body behaves like a tiny magnet, and it can interact with an external magnetic field. The electrons are ___________ by the magnetic field.

Determining the Charge-To-Mass Ratio In 1897, JJ Thomson, an English Physicist, determined the charge-to-mass ratio of an electron by adjusting the electric field so that the _____________ deflection (θE) was the same as the _____________ deflection (θB), and was able to calculate the charge-to-mass ratio of an electron using the following equation:

e / m ratio = EθEB2l

Thomson determined the charge-to-mass ratio of an electron to be -1.76 x 108 coulombs per gram, meaning it was approximately 2000 times lighter than hydrogen, the lightest known atom.

e / m ratio = EθEB2l

= −1.76 ×108 coulombs per gram



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CONCEPT: RUTHERFORD GOLD FOIL EXPERIMENT

The __________________________ experiment also called the Rutherford Gold Foil experiment helped to discover that any given atom had a positively charged center called the _______________.

• It is there where most of the atom’s mass was concentrated.

Subatomic Particle

Charge

Mass

Relative

Absolute

Relative (in amu)

Absolute (in kg)

Proton (p+)

+1

+1.60 x 10-19 C

1.00727

1.673 x 10-27

Neutron (no)

0

0

1.00866

1.673 x 10-27

Electron (e –)

– 1

-1.60 x 10-19 C

5.49 x 10-4

9.11 x 10-31



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CONCEPT: MILLIKAN OIL DROP EXPERIMENT In 1913 Robert Millikan and Harvey Fletcher discovered the charge of an electron as being _____________________ .

The charge of an electron When an oil droplet is suspended, mass x acceleration (m x g) due to gravity is exactly counterbalanced by the electric force applied. The electric force applied equals the applied electric field E times the charge on the drop (q). Making them equal to one another: The mass of an electron By using his discovered charge and then the charge-to-mass ratio determined by Thomson’s cathode ray tube experiment we are able to calculate the mass of electron.



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CONCEPT: CHADWICK NEUTRON EXPERIMENT In 1920, Ernest Rutherford stated that the nucleus must contain neutral, massive particles.

In the early 1930s with experiments designed by Walter Bothe as well as Mr. and Mrs. Joliot it was determined that bombarding ______________ with alpha particles would produce high-energy radiation.

In 1932, James Chadwick modified the earlier experiments and determined that the unknown neutral particles in the nucleus were the _____________.

• By examining the motion of these neutral and unknown particles, Chadwick was able to determine the velocity of the protons.

• Through ________________________________________ he determined that the mass of the neutral particles were nearly identical to the mass of a proton.

Relative (in amu)

Absolute (in kg)

Proton (p+)

1.00727

1.673 x 10-27

Neutron (no)

1.00866

1.673 x 10-27

• His equation to prove the existence of this neutral particle can be written as:



Page 15

CONCEPT: ATOMIC PROPERTIES AND CHEMICAL BONDS

Before we examine the types of chemical bonding, we should ask why atoms bond at all.

• Generally, the reason is that ionic bonding ____________ the potential energy between positive and negative ions.

• Generally, the reason covalent bonds form is to follow the ____________ rule, in which the element is then

surrounded by 8 valence electrons.

There are three models of chemical bonding:

In ____________________ bonding, metals connect to non-metals.

• __________________ transfers an electron to the ________________ , creating ions with opposite charges that

are attracted to each other.

Li F Li F Li F

In _______________ bonding, non-metals connect to non-metals.

• In it the nonmetals __________________ electron pairs between their nuclei.

ClCl

In _______________ bonding, metal atoms “pool” their valence electrons to form an electron “sea” that holds the metal-ion

together



Page 16

CONCEPT: CHEMICAL BONDS (PRACTICE)

EXAMPLE: Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound.

atomic element ––

molecular element ––

molecular compound ––

ionic compound ––

a. Iodine

b. NH3

c. Graphite

d. Na3P

e. Ag2(SO4)2



Page 17

CONCEPT: THE IONIC-BONDING MODEL

The central idea of ionic bonding is that the metal transfers an electron(s) to a nonmetal.

• The metal then becomes a(n) ____________ (positive ion). and the nonmetal becomes a(n) _____________ (negative ion).

• Their opposite charges cause them to combine into a crystalline solid.

PRACTICE: Determine the molecular formula of the compound formed from each of the following ions.

a. K+ & P3-

b. Sn4+ & O2-

c. Al3+ & CO32-



Page 18

CONCEPT: COMMON POLYATOMIC IONS

Polyatomic ions are compounds made up of different elements, usually only ____________, and possess a ____________.

Singly Charged Cation (Positive Ion)

NH4+

Ammonium

Doubly Charged Cation (Positive Ion)

Hg22+

Mercury (I)

Singly Charged Anions (Negative Ions)

CH3CO2– or C2H3O2

–

Acetate

CN–

Cyanide

OH–

Hydroxide

MnO4–

Permanganate

NO3–

Nitrate

Nitrite

Doubly & Singly Charged Anions (Negative Ions)

HPO42–

Hydrogen Phosphate

H2PO4–

Dihydrogen Phosphate

HCO3–

Hydrogen Carbonate or Bicarbonate

HSO4–

Hydrogen Sulfate or Bisulfate

Doubly Charged Anions (Negative Ions)

CO32–

Carbonate

CrO42–

Chromate

Cr2O72–

Dichromate

O22–

Peroxide

SO42–

Sulfate

Sulfite

Triply Charged Anions (Negative Ions)

PO43–

Phosphate

Phosphite



Page 19

CONCEPT: POLYATOMIC IONS w/ HALOGENS

Polyatomic ions containing halogens are sometimes referred to as __________ halogens or halogen ________________.

These compounds share 4 common characteristics:

1.

2.

3.

4.

These compounds use the same system for naming:

PRACTICE: Name each of the following compounds.

a. BrO4 – b. FO2 –

c. ClO – d. IO3 –



Page 20

CONCEPT: NAMING MOLECULAR COMPOUNDS Features: _________________ & _________________

Because molecular compounds combine in different proportions to form different compounds, we must use numerical

prefixes.

Rules for Naming: a. The first nonmetal is named normally and uses all numerical prefixes except ___________________. b. The second nonmetal keeps its base name but has its ending changed to _____________________. EXAMPLE: Write the formula for each of the following compounds.

a. Disulfur monobromide b. Iodine Tetrachloride

PRACTICE: Give the systematic name for each of the following compounds:

a. CO b. N2S4 c. IO5



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CONCEPT: IONIC COMPOUNDS

In the early days of chemistry, newly discovered compounds were given fancy names such as morphine, quicklime and

muriatic acid. Since then thousands of new compounds have been discovered and named under a system called

_____________________________.

Metals tend to __________ electrons to become positively charged ions called _______________.

Nonmetals tend to __________ electrons to become negatively charged ions called _______________.



Page 22

CONCEPT: NAMING BINARY IONIC COMPOUNDS Features: ___________________ & ___________________

Rules for Naming: a. The metal is named and written first.

• If the metal is a transition metal we must use a _________________________ to describe its positive charge.

b. The nonmetal keeps its base name but has its ending changed to ___________________.

EXAMPLE: Provide the molecular formula or name for each of the following compounds.

a. Calcium phosphide b. CoO

PRACTICE: Provide the molecular formula or name for each of the following compounds.

a. AlBr3 b. Lead (IV) sulfide c. SnO2



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CONCEPT: NAMING IONIC COMPOUNDS w/ POLYATOMICS Features: _________________ & _________________

Rules for Naming:

a) The metal keeps its name and is named and written first.

• If the metal is a transition metal we must use a _____________________ to describe its positive charge.

b) Name the polyatomic as you would normally.

EXAMPLE: Write the formula for each of the following compounds:

a. Iron (III) Acetate b. Copper (I) phosphate

c. Strontium Carbonate d. Ammonium Nitrite

EXAMPLE: Give the systematic name for each of the following compounds:

a. Pb(CrO4)2 b. Ga(ClO4)3

c. Mn(HSO4)2 d. Ba(CN)2



Page 24

CONCEPT: NAMING IONIC HYDRATES Features: _________________ & _________________

CuSO4 5 H2O

Rules for Naming the Ionic Compound portion: a. The metal is named normally and written first.

• If the metal is a transition metal we must use a ________________________ to describe its positive charge.

b. The nonmetal keeps the first part of its name but has its ending changed to ___________________.

c. Name the polyatomic as you would normally.

Rules for Naming the H2O portion:

a. The H2O portion will be called ___________________ .

b. To describe the number of H2O molecules use these prefixes.

EXAMPLE: Write the formula for each of the following compounds.

a. Calcium carbonate hexahydrate

b. Lead (IV) Sulfate pentahydrate


a) K2Cr2O7 · 3 H2O b) Sn(SO3)2 · 4 H2O



Page 25

CONCEPT: NAMING ACIDS

1. BINARY ACIDS Features: _______________________ + _______________________

Rules for Naming:

a. The prefix will be ___________________ .

b. Use the base name of the nonmetal.

c. The suffix will be ___________________ .

EXAMPLE: Write the formula for each of the following compounds:

a. Hydroiodic acid b. Hydroselenic acid c. Hydrofluoric acid


a. HBr b. H2S c. HCN

2. OXOACIDS or OXYACIDS Features: _______________________ + _______________________

Rules for Naming: a. If the polyatomic ion ends with –ate then change the ending to _____________________. b. If the polyatomic ion ends with –ite then change the ending to ______________________.

EXAMPLE: Give the systematic name or formula for each of the following compounds:

a. H2CO3 b. Nitric acid c. H2SO4

PRACTICE: Give the systematic name or formula for each of the following compounds:

a. Hypobromous acid b. HClO3 c. Acetic acid



Page 26

3. How many molecules of hexane are contained in 55.0 mL of hexane? The density of hexane is 0.6548 g/mL and the molar mass is 86.17 g/mol.



Page 27

4. How many SO3 ions are contained in 120.0 mg of Na2SO3? The molar mass of Na2SO3 is 126.05 g/mol.



Page 28

5. What mass of phosphorus pentafluoride, PF5, has the same number of fluorine atoms as 50.0 g of oxygen difluoride, OF2?



Page 29

6. How many bromide ions are there in 4.50 moles of gallium bromide?



Page 30

7. How many moles of oxygen atoms are required to combine with 3.05 moles of Pb to create lead (IV) phosphate?



Page 31

8. How many cations are there in 100.0 g of lithium nitride?



Page 32

10. Which of the following amounts would contain the least atoms? a) 10.0 g Sr

b) 10.0 g Br

c) 10.0 g Mg

d) 10.0 g Li



Page 33

11. Which of the following amounts have the most molecules?

a) 15.0 g N2

b) 15.0 g Br2

c) 15.0 g O2

d) 15.0 g I2



Page 34

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CHEMISTRY - BROWN 13E CH.2 - ATOMS, MOLECULES & IONSlightcat-files.s3.amazonaws.com/packets/admin... · Atoms Moles Grams EXAMPLE: Determine the mass (in grams) found in 7.28 x 1028