The Scope of Chemistry Thinking Like a Scientist - LPSisite.lps.org/mgeist/N-C-U1.pdf · •The Scope of Chemistry ... • The study of the composition of matter and the changes that

Chapter 1: Introduction to Chemistry

• The Scope of Chemistry• Thinking Like a Scientist


The Scope of Chemistry

Chapter 1: Introduction to Chemistry-- The Scope of Chemistry --

Chemistry• The study of the composition of matter and the changes that matter

undergoes• Major areas of chemistry

– Organic chemistry: study of essentially all substances containing carbon

– Inorganic chemistry: study of essentially all substances NOT containing carbon

– Analytical chemistry: study of the composition of substances (by percent, quantity, etc.)

– Physical chemistry: study of chemistry from a standpoint of mathematics and physics

– Biochemistry: study of the chemistry of living organisms


Thinking Like A Scientist

Chapter 1: Introduction to Chemistry-- Thinking Like a Scientist --

Scientific method• A method of inquiry involving observations, experiments, hypotheses,

and broad explanations called theories• Used in biology, chemistry, physics, finance, etc.Steps of the scientific method Observation

– Information obtained through the senses– Often involves a measurement

Hypothesis– A proposed explanation for observations– NOT the same as a scientific theory or law– Often referred to as “an educated guess”

Experiment– A carefully-controlled, repeatable procedure for gathering data in

order to test a hypothesis– Can either support or discard a hypothesis– Experimental flaws/errors

Chapter 1: Introduction to Chemistry-- Thinking Like a Scientist --

Data Analysis– Analyzing data that has been recorded during the experiment– Looking for trends in data to support or discard the hypothesis

Conclusions– Results from analyzing data and checking for errors in the

experiment– Serves as a summary to support or discard the hypothesis

Scientific theory Scientific lawBroad and extensively-tested explanation of why experiments yield certain results

Concise statement that summarizes the results of many observations and experiments

Can never be proven but can be disproven

Widely-accepted

Explains why things happen Explain what things happen

Chapter 2: Matter and Change• Properties of Matter• Mixtures• Elements and Compounds• Chemical Reactions

Chapter 2: Matter and Change

Properties of Matter

Chapter 2: Matter and Change-- Properties of Matter --

Matter• Anything that takes up space and has mass• Substance

– Has uniform and definite composition– Contains only one kind of matter

• Properties– Physical properties

• Qualities of substances that can be observed or measured without changing the substance’s chemical composition

• Examples– Solubility– Odor– Hardness– Density– Melting point– Boiling point


– Chemical properties• Abilities of substances to undergo chemical reactions and to

form new substances• Examples

– Flammability– Reactivity to oxygen– Reacts with acids

– Physical states of matter• Determine many distinguishing characteristics of matter• Deal with the same substance but in different physical means of

existence


Physical change• A change in the form of a substance, not its chemical composition• Examples

– Boiling water– Freezing mercury– Cutting paper– Dissolving salt in water

Chemical change• A change in which a given substance becomes a new substance or

substances with different properties and different composition• Examples

– Burning paper– Detonating an explosive– Grilling a steak– Reacting an acid with a base

Chapter 2: Matter and Change-- Properties of Matter --Physical States of Matter

States of MatterProperty Solid Liquid Gas

Shape Definite shape Shape depends on container

Shape depends on container

Volume Definite volume Definite volume Indefinite volume

Expansion/Compressibility

Some expansion but minimal

Some expansion when heated

Easily compressed and

expandedExample with

Water Ice Water Steam

Specific notes

Gas: gaseous state of a substance not generally aliquid or solid at room temperature

Vapor: gaseous state of a substance that isgenerally a liquid or solid at roomtemperature


Law of conservation of mass• Scientific law that states that mass can be neither created nor

destroyed in an ordinary chemical or physical process• The total mass of products = The total mass of reactants• Includes all gases, liquids, and solids involved in the chemical reaction


Mixtures

Chapter 2: Matter and Change-- Mixtures --

Mixture• Physical blend of two or more substances• Compositions may vary.• Classifications

– Heterogeneous mixture• A mixture that is not uniform in composition• Example: salad• Portions differ from other portions

– Homogeneous mixture• A mixture that is uniform in composition• Example: saltwater• Portions the same as other portions• Solution

– Special name of a homogeneous mixture– Can be same or different phases of matter together

Chapter 2: Matter and Change-- Mixtures --

– Kinds of solutions» Gas-gas: carbon dioxide and oxygen (air)» Liquid-gas: water vapor in air (moist air)» Gas-liquid: oxygen in water (ocean water)» Liquid-liquid: hydrochloric acid in water» Solid-liquid: saltwater (salt in water)» Solid-solid: copper in silver (sterling silver)

• Separating mixtures– Distillation– Magnetism– Filters– Other methods


Elements and Compounds

Chapter 2: Matter and Change-- Elements and Compounds --

Element• Simplest form of matter that can exist under normal laboratory

conditions• Can be found on the Periodic Table of Elements• Building blocks for all other substances• Atoms are each classified by names of elements• Examples

– Carbon (C)– Oxygen (O)– Nitrogen (N)– Sodium (Na)– Lead (Pb)– Copper (Cu)– Cobalt (Co)


Compound• Substance that can be separated into simpler substances only by

chemical means• Two or more elements that chemically combine with one another• Examples

– Carbon monoxide (CO)– Carbon dioxide (CO2)– Sodium chloride (NaCl)– Glucose (C6H12O6)– RDX (C3H6N6O6)– Acetone (CH3COCH3)



Symbol• One- or two-letter representation for an element• Letters for symbol are generally English, Latin, or Greek in origin• Examples

– Na (Latin from natrium for “sodium”)– Au (Latin from aurum for “gold”)– He (English for “helium”)– W (German from wolfram for “tungsten”)– Pb (Greek from plumbum for “lead”)


Chemical Reactions

Chapter 2: Matter and Change-- Chemical Reactions --

Chemical reaction• The changing of substances to other substances by the breaking of

bonds in the reactants and the formation of bonds in the products• Reactants

– Starting substances in a chemical reaction– On the left side of a chemical reaction equation

• Products– Ending substances in a chemical reaction– On the right side of a chemical reaction equation

• ExampleHydrochloric acid + Sodium hydroxide Sodium chloride + Water

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

Chapter 3:Scientific Measurement

• The Importance of Measurement

• Uncertainty in Measurements• International System of Units• Density• Temperature


The Importance of Measurement

Chapter 3: Scientific Measurement-- The Importance of Measurement --

• Qualitative and quantitative measurements– Use of the International System of Measurements (SI)– Qualitative measurements

• Measurements that give descriptive, nonnumeric results• Examples: color, touch, etc.• Vary in measurement from person to person

– Quantitative measurements• Measurements that give definite, usually numeric results• Examples: temperature, speed, acceleration, distance, time• Can be no more reliable that the instrument used to make the

measurement and the care with which it is used and read• Scientific notation

– Expression of numbers in the form a x 10b where a is equal to or greater than 1 and less than 10 and b is an integer

– Used to abbreviate very big or very small numbers


– Examples• If the exponent of 10 is positive n, move the decimal point n

spaces to the right and fill in extra spaces with zeros.• If the exponent of 10 is negative n, move the decimal point n

spaces to the left and fill in extra spaces with zeros.– Exponents on 10 must be the same to add and subtract.




Uncertainty in Measurements

Chapter 3: Scientific Measurement-- Uncertainty in Measurements --

• Accuracy vs. Precision– Accuracy

• The closeness of a measurement to the true value of what is being measured

• Example: data of 35, 36, and 35 when the true value is 34.5– Precision

• The closeness or reproducibility of a set of measurements taken under the same conditions

• Examples– 35, 36, and 35 in a trial where the true value is 48.5– 35, 36, and 35 in a trial where the true value is 34.5

• Error analysis– Accepted value: correct value based on reliable references– Experimental value: value measured in the laboratory


– Calculating error and percent errorError = |accepted value – experimental value|

– ExampleA scientist is doing an experiment in which he measures the density of water. He experimentally finds the average of the densities found to be 0.982 g/mL. The accepted value is 1.000 g/mL. What is his error and percent error?


• Significant figures– All the digits that can be known precisely in a measurement, plus a

last estimated digit– Example: a graduated cylinder that can be read to the centiliter

must be estimated to the milliliter (read: 0.7 mL significant: 0.74)– Must know the equipment to know where it is exact and where you

need to guess– Rules for significant figures

• Every nonzero digit in a reported measurement is assumed to be significant. (2.853 has 4 significant digits)

• Zeros appearing between digits are significant. (10.002 has 5 significant digits)

• Leftmost zeros appearing in front of nonzero digits are not significant. (0.000032 has 2 significant digits)

• Zeros at the end of a number and to the right of the decimal point are always significant. (34.000 has 5 significant digits)


• Zeros at the rightmost end of a measurement that lie to the left of an understood decimal point are not significant. (340000 has 2 significant digits)

• Two cases of unlimited number of significant digits– Direct counting– Exactly defined quantities for converting units


Problem Set• 1.2543• 10.543• 9870.0• 9870.• 0.000025• 0.00000250• 78000006• 20 students• 1000 mL = 1 L

5 significant figures5 significant figures5 significant figures4 significant figures2 significant figures3 significant figures8 significant figures significant figures significant figures


• Rules for where to round– The answer to an addition or subtraction calculation should be

rounded to the same number of decimal places as the measurement with the least number of decimal places.

Example: 39.892 + 9.45 + 6.9075 = 56.2495Since the least accurate measurement is 9.45 and it has only two decimal places, the answer must be rounded to two decimal places.Answer: 56.25


– The answer to a multiplication or division calculation should be rounded to the same number of significant digits as the measurement with the least significant digits.

Example: 7.546 x 2.58 x 18.325 = 356.763561Since the measurement with the least significant digits is 2.58 and it has only three significant digits, the answer must be rounded to three significant digits.Answer: 357

• International System of Units (SI System)– Universal system of units– Seven SI units to know

• Length: meter (m)• Mass: kilogram (kg)• Temperature: kelvin (K)• Time: second (s)• Pressure: pascal (Pa)• Energy: joule (J)• Amount of substance: mole (mol)

Chapter 3: Scientific Measurement-- International System of Units --

• Prefixes to use that represent relationships (three to know)

Chapter 3: Scientific Measurement-- International System of Units --

Centi- (c)• 100 cm = 1 m• 100 cg = 1 g• 100 cGeist = 1 Geist

Milli- (m)• 1000 mm = 1 m• 1000 mL = 1 L• 1000 mGabriel = 1 Gabriel

Kilo- (k)• 1 km = 1000 m• 1 kg = 1000 g• 1 kAnthony = 1000 Anthony

• Converting units and measurements– 5 1 = 5– Conversion factor

• Ratio of equivalent measurements• The same as multiplying times one (1)• Example

– 100 cm = 1 m–

Chapter 4: Problem Solving in Chemistry-- Conversion Problems --

– When converting to other units, start with what you are given, and use conversion factors to change to cancel out units to get what you want.

• Dimensional analysis (basic conversions)– How many centimeters are in 5.2 kilometers?

• 1 km = 1000 m• 1 m = 100 cm


(2 significant figures)= 520000 cm (or 5.2 105 cm)

2 s.f. s.f. s.f.

• Dimensional analysis (basic conversions)– Convert 40.0 km/hr to m/s.

• 1 km = 1000 m• 60 s = 1 min• 60 min = 1 hr


(3 significant figures)= 11.1 m/s

3 s.f. s.f. s.f. s.f.

• Dimensional analysis (advanced conversions)– The speed limit in residential zones is 25 miles per hour.

If you were living in the UK, what would this speed limit be in meters per second?

• 5280 ft = 1 mi• 12 in = 1 ft• 2.54 cm = 1 in (back of periodic table)• 1 m = 100 cm• 60 min = 1 hr• 60 s = 1 min



(2 significant figures)= 11 m/s

2 s.f. s.f. s.f. s.f. s.f. s.f. s.f. • Dimensional analysis (advanced conversions)– Tire pressure in your car tires should be 35.0 psi

(pounds per square inch or lb/in2). What pressure is this in kilograms per square centimeter (kg/cm2)?

• 1 lb = 0.4536 kg (back of periodic table)• 1 in = 2.54 cm (back of periodic table)



(3 significant figures)= 2.46 kg/cm2

3 s.f. s.f. s.f.

Chapter 3: Scientific Measurement-- Density --

Density• Ratio between the mass and the volume of a substance• Calculated by taking the mass and dividing by the volume

• Density is a rate, meaning it can be calculated from a slope.

• Independent data on the horizontal axis (x-axis) -- volume• Dependent data on the vertical axis (y-axis) -- mass


y = mx + b (Slope-intercept form)m = dV + b

y = mass (on y-axis)d = density (slope)x = volume (on x-axis)


Write the equation of the line for substance A.

y = mx + b

m = 0.789b = –0.00145

m = 0.789V – 0.00145


Calculate the mass of a 14.0 cm3 piece of substance A.

1 cm3 = 1 mL, so 14.0 cm3 = 14.0 mL.

m = 0.789V – 0.00145m = 0.789(14.0) – 0.00145m = 11.0 – 0.00145 (least number of significant figures due to x/÷)m = 11.0 g (least number of decimal places due to +/–)


The density of substance A is 0.789 g/mL.

Based on the graph and the table, what is substance A?

Substance A is ethanol.


What is the volume (in mL) of an object with a density of 7.02 g/cm3 and a mass of 6.00 x 102 g?

Density formula

Setting fraction equal to fraction so that you can cross-multiply

Cross-multiplying

Getting V by itself

Substituting values (keeping in mind cm3 = mL)

Dividing

Documents

The Scope of Chemistry Thinking Like a Scientist - LPSisite.lps.org/mgeist/N-C-U1.pdf · •The Scope of Chemistry ... • The study of the composition of matter and the changes that