Introduction to Chemistry 1Chapter 1 Overview   An understanding of the history of chemical investigation.

• The history of experimentation and scientific inquiry.

1.1-1.2 The Scientific Method: How Chemists Think

Observation – hypothesis – law – theory – experimentScientific law (e.g. law of conservation of massDalton's atomic theory)

REVIEW CHAPTERS 1, 2, 3 and 4

1.3 Matter What Is Matter? A. Occupies space and has mass B. Atom – smallest unit of matter C. Molecule – atoms joined together

1.4 Classifying Matter According to Its State: Solid, Liquid, and Gas A. Solid (fixed volume, incompressible) strongest attractive forces

1. Crystalline2. Amorphous

B. Liquid- Medium inter molecular forces of attraction1. Fixed volume2. Fluid

C. Gas (lot if empty space)1. Compressible2. Fluid3. No attractive forces

Classifying Matter Section 1.4

Numerical Side of Chemistry 2Chapter Overview   A cornerstone of the chemical sciences, the manipulation of numbers and their associated units. Measurement accuracies, significant figures, rounding and scientific notation.

2.1 Scientific Notation: Writing Big and Small NumbersA. Shorthand notation for numbersB. Two main pieces: decimal and power-of-10 exponentC. Measured value does not change, just how you report it (550.6 to 1 sig fig?)

2.2, 2.4 Numbers in chemistry–Units and precision and accuracy in reporting it; Uncertainty in measurement, last digit on the right is uncertain

2.3 Significant Figures: Writing Numbers to Reflect PositionA. How many digits can I report? How many should I report?B. Certain digits and estimated digitsC. Counting significant figures

1. All nonzero digits are significant 1234 = 4 Sig fig 2. Interior zeros are significant 505 = 3 sig fig

3. Trailing zeros after a decimal are significant 55.00 = 4 sig fig4. Leading zeros are not significant 0.012 = 2 sig fig5. Zeros at the end of a number, without a decimal point, are ambiguous 150 = 2 SF

D. Exact numbers/definition, conversion factors have infinite sig fig.

2.4 Significant Figures in Calculations

A. Multiplication and division:

Result carries as many significant digits as the factor with the fewest significant digits

B. Rounding

1. If leftmost dropped digit is 4 or less, round down (leave it same)

2. If leftmost dropped digit is 5 or higher, round up (increment it by 1)

C. Addition and Subtraction

Result carries as many decimal places as the quantity with the fewest decimal places

D. Calculations Involving Both Multiplication/Division and Addition/Subtraction

1. Do steps in parentheses first2. Determine the number of significant figures in intermediate answer3. Do remaining steps

2.5 The Basic Units of MeasurementA. English, metric, SIB. SI Units (Mass – kg; Length – m; Time – sec)C. Prefix Multipliers

milli (m) 0.001 centi (c) 0.01 kilo (k) 1000 Mega (M) 1,000,000

D. Derived Units1. Area – cm2

2. Volume – cm3 or L

Common Prefixes in the SI System MEMORIZE

Prefix SymbolDecimal

EquivalentPower of 10

mega- M 1,000,000 Base x 106

kilo- k 1,000 Base x 103

deci- d 0.1 Base x 10-1

centi- c 0.01 Base x 10-2

milli- m 0.001 Base x 10-3

micro- or mc 0.000 001 Base x 10-6

nano- n 0.000 000 001 Base x 10-9

2.6 Converting from One Unit to Another (UNIT 1 to UNIT 2)A. Units are important, most numbers get oneB. Include units in all calculationsC. Conversion factors (Unit you have comes in the bottom, unit you want comes in the TOP)

Unit 1 X Unit 2 = UNIT 2 Unit 1

D. Significant figure of the final answer depends on UNIT 1 given in problem NOT the sig. fig of the conversion factor

Solving Multistep Conversion ProblemsA. Understand where you are going firstB. Not all calculations can be done in one step

Units Raised to a PowerA. 1 inch = 2.54 cm so 1 inch3 = (2.54 cm)3 = 16.4 cm3

2.7 Temperature: Random Molecular and Atomic Motion

A. Fahrenheit (F) B. Celsius (C)C. Kelvin (K)

Conversions

2.8 Density = Mass/Volume; Remember Mass is in grams, volume in mL or cm3

Unit of density = g/mL or g/cm3

1.8

32-F C

273C K

F 23C)(1.8

Atoms and Elements 33.1 ELEMENTS A. Atoms make up all matter

B. ~91 different naturally occurring elements 1. Each element is composed of tiny indestructible

particles called atoms 2. All atoms of a given element have the same

mass and other properties that distinguish them from atoms of other elements

3. Atoms combine in simple, whole-number ratios to form compounds3.3, 3.4 Names vs symbols of elements- Memorize List from webMetal, non-metals, metalloids-

3.5 and 3.6 Looking for Patterns: The Periodic Law and the Periodic Table

A. Mendeleev (1834 - 1907)B. Periodic lawC. Metals (usually loose electrons to form cations)D. Nonmetals (usually gains electrons to form anions)E. Metalloids, also known as semiconductorsF. Individual group names

1. Group 1 – alkali metals +1 ion only2. Group 2 – alkali earth metals +2 ion only3. Group 7 – halogens usually –1 ion4. Group 8 – noble gases

3.7 Elements which are diatomics N2 O2 F2 Cl2 Br2 I2 and H2

3.8 Ionic compound – It must conatin a metal, Meatl forms + ion CATIONSNonmatels form – charge ions ANIONS

3.9 Number of atoms in a chemical formula

Section 4.1-4.2How We Tell Matter Apart: Physical and Chemical Properties

A. Physical property1. Observable without changing the identity2. Melting point, odor, color

B. Chemical property1. Observable only by changing the identity-Chemical reactions2. Flammability

How Matter Changes: Physical and Chemical ChangesA. Physical change

1. Appearance and properties can change2. Composition does not change

B. Chemical change1. Appearance and properties can change2. Composition changes

Section 4.3 Separation of mixtures through physical changes

1. Decanting2. Distillation3. Filtration

4 Properties of Matter

Section 4.4 EnergyA. Energy cannot be created or destroyed

B. Units of energy and heat

1. Joule (J)

2. calorie (cal) (1 cal = 4.184J)

3. Calorie (Cal) (1Cal = 1000cal = 1kcal)

4. Kilowatt-hour (kWh)- - Will not be used in CHE100

Exothermic Process: Heat is Released by a system or process

Heat q = negative; Here products are lower energy than reactants

Endothermic Process: Heat is Absorbed by a system or process

Heat q = positive; Here products are HIGHER energy than reactants

Section 4.5 Heat, Specific heat

The specific heat of a substance is the quantity of heat required to change the temperature of 1 g of that substance by 1ºC units J/gºC

Specific heat of water very high

Specific heat of metals are low, hence are good conductors of heat

Heat = mass of substance X Sp. Heat X Change in temp

= M S T

2 body problem Heat lost by hot object = heat gained by cold object

Read definitions of Specific heat (cal/gºC) meaning of itSpecific of heat of water = 1cal/g ºC = 4.184J/g ºC

EXAM 1- 100 POINTS10 points Bonus question!!

Part 1 Multiple Choice –Show calculations for partial/full credit20 questions- chapters (1-4)

Part 2 Fill in the Blanks scientific notation, chemical change, physical change, chemical and physical properties, Pure substances and mixtures, Simple Conversions –show all work and Density, calories problem

Bonus question: