Ch. 1 Introduction: Some Basic Concepts

Chemistry

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

Matter

• Anything that has mass and takes up space

States of Matter

Solid

• Definite shape AND volume• Incompressible• Particles are packed tightly together

Liquid

• Definite volume• Takes the shape of its container• Almost incompressible-particles not rigidly

packed

Gas• Takes the volume and the shape of its

container• Particles in a gas are spaced far apart• Easily compressed

Vapor

• The gaseous state of a substance that is generally a solid or a liquid at room temperature

Physical Properties• A quality or condition that can be observed or

measured without changing the substances composition

• Color, solubility, odor, hardness, density, melting point, boiling point

Physical Change

• Changes the material without changing the composition

• Boiling, freezing, dissolving, melting, condensing, breaking, splitting, cracking, cutting, crushing, bending………

• Usually reversible

Pure Substance

• Contain only one kind of matter• Have identical physical properties

Elements

• The simplest forms of matter that can exist under laboratory condition

• Can not be separated into simpler substances by chemical means

• The building blocks for all other substances

Chemical Symbol

• One or 2 letters• The first letter is always capitalized

Compounds

• 2 or more elements chemically combined• Can be separated into simpler substances by

chemical means

Sodium Metal

+Chlorine Gas

=Sodium Chloride

Compounds• The subscript numbers in chemical formulas

represent the proportions of elements that make up the compounds

• Pb2(SO4)3

• Pb (lead) 2• S(Sulfur) 3• Oxygen 12

Law of Definite Proportions (Law of Constant Composition)

• The elemental composition of a pure compound is ALWAYS the same

• i.e. Water= H2O = 2 Hydrogens: 1 Oxygen ALWAYS

2 Types of Mixtures

• Heterogeneous Mixtures• Homogeneous Mixtures

Mixture

• A physical blend of 2 substances• Compositions may vary

Heterogeneous Mixture

• One that is NOT uniform in composition

Homogeneous Mixture

• The same throughout

Solution

• A homogeneous mixture• Solutions may be solids, liquids, or gases• Same composition throughout

Phase

• Any part of a system with uniform composition and properties

Phase 1

Phase 2

Separation of Mixtures

Distillation

Separation of Mixtures

Centrifuge

Separation of Mixtures

Chemical Property• The ability of a substance to undergo a

chemical reaction to form a new substance

• I.e. flammability, alkalinity, acidity, rusting

• Chemical properties are only observed when a substance undergoes a chemical change

Chemical Changes (Reaction)

• One or more substances change into new substances

• 2 H2 (g) + O2 (g) 2 H2O (g)

Reactants Products

Chemical Reactions (rust-oxidation)

Chemical Reactions (combustion)

Chemical Reactions (acid/base)

Indicators of a Chemical Reaction

1. Energy is absorbed or given off (change in temperature)

2. Change in color3. Change in odor4. Formation of a solid (precipitation)5. Formation of a gas

Scientific Method

• One logical, systematic approach to the solution of scientific problems. Steps include:

1. Making observations2. Testing hypothesis3. Developing theories

Observation

• Use your senses to obtain information directly

Hypothesis

• A proposed explanation for an observation based on previous knowledge (or research)

• Must be specific• Must be testable• Is only useful if it accounts for what is actually

observed

Experiment

• A means to test a hypothesis

Manipulated Variable (Independent Variable)

• The variable that you can change• Time• Temperature• Volume• Speed• Pressure

Independent Variable

Responding Variable (Dependent Variable)

• The variable that is observed during the experiment

Dependent Variable

• For the results of an experiment to be accepted the experiment must produce the same results no matter how many times it is repeated or by whom

Theory

• A broad and extensively tested explanation of why experiments give certain results.

• A theory can NEVER be proven because a new experiment can always disprove it

Scientific Law

• A concise statement that summarizes the results of many observations and experiments.

• Scientific law describes natural phenomena without attempting to explain it

Qualitative Observation

• Give results in a descriptive non-numerical form

• Subjective• Ex. The solution is green, The precipitate is

fluffy

Quantitative Observation

• Gives results in a definite form• Numbers and units

Measurement

• A quantity that has both a number and a unit

• Measurements are only as exact as the instrument used to take it

International System of Units (SI)

• Length = meter(m)• Mass = gram (g)• Temperature = kelvin (K) although often we

will use Celsius• Time = second (s)• Amount of substance = mole (mol)• Volume = liter (L)

Prefixes

• Mega (M) = 106

• Kilo (k) = 103

• Deci (d) = 10-1

• Centi (c) = 10-2

• Milli (m) = 10-3

• Micro () 10-6

• Nano (n) = 10-9

• Pico (p) = 10-12

Length

• The distance between 2 points• Unit: meter (m)

Mass

• The amount of matter an object has• Units: Grams• Measuring tools: Triple Beam Balance,

Electronic scale, Analytical Scale

Temperature

• Measure of how hot or cold an object is• Determines heat transfer (moves from high to

low)• Almost all substances expand when heated-

contract when cooled (except water)• Units: Kelvin (K) = Celsius (c) + 273

Volume

• The amount of space an object (substance) occupies

• Units: Liter (L)

Density

• The amount of matter in a given volume

• Density = mass volume• D= m/v

Density

• Depends only on the composition of a substance NOT on the size of the sample

• With a mixture the density can vary because the composition of the mixture can vary

Density

• The density of a substance generally decreases as its temperature increases therefore temperature of the substance must always be noted

Accuracy

• how close a measurement comes to the actual or true value of whatever is measured must be compared to the correct value

Precision

• How close a series of measurements are to one another depends on more than one measurement

Significant Figures

• In a measurement, includes ALL of the digits that are known plus the last digit that is estimated

• Measurements MUST ALWAYS be reported to the correct number of sigfigs because calculated answers depend on sigfigs

Rules for Significant Figures

1. Nonzero numbers are ALWAYS significant-13.24 cm = 4 sigfigs- 3.5 mL = 2 sigfigs-123.456 g = 6 sigfigs

Rules for Significant Figures

2. Captive zeros are ALWAYS significant- 1023 g = 4 sigfigs- 1.0005 cm = 5 sigfigs

Rules for Significant Figures

3. Leading Zeros (zeros in front of nonzero numbers) NEVER count- 0.53 cm = 2 sigfigs- 0.0000000000053 cm = 2 sigfigs

Rules for Significant Figures

4. Zeros at the end of a number only count if there is a decimal point- 100 mL = 1 sigfig-100.0 mL = 4 sigfigs-100.00 mL = 5 sigfigs

Rules for Significant Figures

5. Scientific Notation: Only consider the coefficients when determining sigfigs- 1.2 x 104 mL = 2 sigfigs- 1.06 x 10-6 m = 3 sigfigs

Rules for Significant Figures

6. Exact numbers (counting) or exactly defined quantities have unlimited number of sigfigs- 35 students = unlimited sigfigs- 1 min = 60 seconds = unlimited sigfigs

Significant Figures in Calculations

• A calculated value can NEVER be more precise than a measured value

Significant Figures in Calculations

• Addition and subtraction:• The answer should be rounded to the same number of

decimal places as the measurement with the least number of decimal places

• 135.75 mL (2 decimal places)+ 57.6 mL (1 decimal place) 193.35 mL = 193.4 mL (answer has 1 decimal place)

Rounding

• X > 5 Round up

• X < 5 Round down

• Only look at the number next to the one you are rounding

Significant Figures in Calculations

• Multiplication and Division• Round the answer to the same number of significant

figures as the number with the least number of sigfigs.

• 35.062 g ( 5 sigfigs) = 1.492 g/mL (4 sigfigs) 23.50 mL (4 sigfigs)

Dimensional Analysis

1. Always start with the number given in the problem)

2. Units that you want to cancel go on the bottom

3. When a number is on top, multiply4. When the number is on the bottom, divide

Example

• Convert 153 lb to grams (hint: 1 lb = 453.6 g)

Example

• 153 lb X 453.6 g = 69400.8g=69400g (3sigfigs) 1 lb