1. Unit 1: Sigfigs and Nomenclature
AP Chem
Mrs.Grein

2. Matter A. Pure Substance: doesn’t change composition, fixed properties
1. Molecule: covalent bond
2. Compound: ionic bond
B. Mixture: two or more pure substances
1. Properties change

3. C. States of Matter
1. Solid: strong attraction, low kinetic Energy, constant volume and shape
2. Liquid: constant volume, shape changes
3. Gas: no set volume or shape

4. Energy
A. Definition: ability to do work, apply force to make molecules move
1. Kinetic: energy of motion
2. Potential: stored energy due to position
3. ET: total energy is the sum of kinetic and potential
B. 1st Law of Thermodynamics
1. Conservation of energy, energy cannot be created or destroyed

5. II. Significant Figures
A. Accuracy vs. Precision
1. Consistency leads to the correct outcome- precision leads to accuracy
B. Measurements
1. Two components: value and unit

6. 2. Uncertainty: estimate
a.) The last value in any measurement is a guess
b.) all non-zero digits are significant
c.) zeroes that place the decimal are not significant
1.) sig figs
2.) – 3 sig figs
3.) – 1 sig fig
4.) – 8 sig figs

7. d.) If decimal is present, start from the left and move right to first
non-zero, all that follow are sig figs
e.) If no decimal start right to left to first non-zero, all that follow are sig figs

8. Rounding rules for operation
1. Addition and Subtraction
14.36 cm cm = = cm
use least number of decimal places, least precise
2. Multiplication and Division
14.36 cm ÷ cm = = cm
use least number of sig figs

9. 3.) If the number next to the place value to be rounded is a 5, round your answer to an even number
= 6.682
= 6.684

10. IV. Metric System A. Standard units
Mass ~ kg, g Temperature ~ K Energy ~ Joules
Length ~ m Pressure ~ Pa
Volume~ L Amount of matter ~ mol
Current ~ amps Charge ~ coulomb (c)

11. B. Prefixes
Mega (M) milli (m) 10-3
kilo (k) micro (μ) 10-6
deci (d) nano (n) 10-9
centi (c) pico (p) 10-12

12. V. Conversions
A. Using multiplication of unit factors to cancel unwanted units
1. 1 mile = 5280 ft ft
1 mile

13. B. Sample
1. How many millimeters are equal to 1.85 x 10-4 pm?
2. What is the volume of 62.3g of a gas with a density of 1.2 g/L?

14. VI. Law of Conservation of Mass
A. The total mass of substances does not change during a chemical reaction.
Mass Reactant Mass Reactant = Mass Product

15. 1. Calculating the mass of an element in a compound
a. How much N is in 4.55kg of NH4NO3?
The Formula Mass is:
4 x H = 4 x = amu
2 x N = 2 x = amu
3 x O = 3 x = amu
amu
b. Mass ratio is : amu N / amu NH4NO3 =
(It doesn’t matter how much of the sample you have because the ratio never changes.)

16. B. The Law of Constant Composition
1. No matter what its source, a particular chemical compound is composed of the same elements in the same parts (fractions) by mass.
a. The number of atoms does not equal the amount of contributed mass

17. VI. Law of Multiple Proportions
A. If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers.
1. Ex. Nitrogen Oxides 1 and 2
Oxide %N %O
Oxide %N %O

18. 2. Assume you have 100g of each compound (drop %)
Oxide 1. = g O/ g N = 53.32/46.68 = 1.142:1
Oxide 2 = g O/ g N = 69.55/30.45 = 2.284:1
2.284/1.142 = 2:1
Oxide 1 = NO
Oxide 2 = NO2

19. VII. Dalton’s Atomic Theory
A. all matter consists of atoms
B. atoms of one element cannot be converted into atoms of another element
C. atoms of an element are identical in mass and other properties and are different from atoms of any other element
D. compounds result from the chemical combination of a specific ratio of atoms of different elements.

20. E. Today:
1. the atom is the smallest body that retains the identity of an element.
2. elements can only be converted into other elements in nuclear
reactions
3. Isotopes of an element differ in the number of neutrons and thus in
mass number. A sample of an element is treated as though its atoms
have an average mass.
4. Evidence comes from work of many scientists
a. J.J. Thomson: Raisin bun model
b. Ernest Rutherford: Nuclear Model from gold foil experiment
c. Millikan: Oil drop experiment determined charge/mass ratio of an electron

21. F. Isotopic Symbols A = mass (protons and neutrons)
Z = atomic number (protons)
X = atomic symbol

22. G. Average Atomic Mass
1. Measures the mass and relative abundance of an isotope.
2. Uses this formula
Avg. Atomic Mass =
(isotope1avg.mass x % abundance) + (isotope2avg.mass x %abundance)
100

23. 3. What is the percent abundance of two isotopes of boron if the mass of isotope one is amu and isotope two weighs amu

24. IX. Naming and Writing Formulas
A. Non-metals
1. replace ending of non-metal with …ide
B. Metals
1. All metal names remain intact
2. Transition Metals get Roman Numerals for their different ions

25. Element Ion Formula Systematic name Common name
Copper Cu+1 copper (I) cuprous
Cu+2 copper (II) cupric
Mercury Hg+1 mercury (I) mercurous
Hg+2 mercury (II) mercuric
Cobalt Co+2 cobalt (II)
Co+3 cobalt (III)
Iron Fe+2 iron (II) ferrous
Fe+3 iron (III) ferric
Manganese
Mn+2 manganese (II) manganous
Mn+3 manganese (III)
Tin Sn+2 tin (II) stannous
Sn+4 tin (IV) stannic
Lead Pb+2 lead (II) plumbous
Pb+4 lead (IV) plumbic
**( -ic ending is higher charge than –ous ending)

26. C. Naming Binary Compounds
a. Magnesium and Nitrogen:
Mg3N2 ____________________
b. Iodine and Cadmium:
CdI2 _______________________

27. D. Naming and Formulas of Ionic Compounds
a. tin (II) fluoride (stannous fluoride) : SnF2
b. CrI3: chromium (III) iodide (chromic iodide)
c. Ferric Oxide: Fe2O3
d. CoS: cobalt (II) sulfide (cobaltous sulfide)

28. E. Naming Polyatomic Oxyions
ClO per root ate ClO- hypo root ite
ClO3- root ate
ClO2- root ite

29. F. Acid Naming: 1. Based on name of anion. a. binary anion: …..ide
1. acid becomes hydro…..ic acid
a) HCl:
b. oxyion acids
1. …..ate ions become ……ic acids
a) H2CO3:
2. …..ite ions become …..ous acids
a) H3PO3:

30. G. Naming Hydrates and Binary Covalent Compounds
1. Prefixes
1- mono 6- hexa
2- di 7- hepta
3- tri 8- octa
4- tetra 9- nona
5- penta deca

31. H. Alkanes CnH(2n+2)
Saturated hydrocarbons--each C is bonded to the maximum number of other atoms (4)
Naming alkanes: root + suffix (the ROOT tells the # of carbon atoms):
Meth-1 Hex-6
Eth-2 Hept-7
Prop-3 Oct-8
But-4 Non-9
Pent-5 Dec-10

32. Alkenes: CnH2n 1. Contains at least one C=C bond
2. Unsaturated hydrocarbons

33. J. Alkynes: CnH2n-2
Contains at least one triple C bond

34. K. Examples C6H12 C4H6 Octane Nonene Stannic dichromate tetrahydrate
Cupric Acetate

35. L. Formula Types 1. Empirical: lowest possible ratio of atoms
2. Molecular: actual ratio multiple of empirical
3. Structural: shows arrangement
a. Used mainly for hydrocarbons
B. Categories of hydrocarbons
Alkanes: chain of single bonded carbons
Alkenes: chain of carbons with a double bond
Benzene: cyclohexene
3. Alkynes: chain with a triple bond

36. c. Prefixes for chains and branches
1) if dealing with a chain add: -ane ; -ene ; -yne
2) if dealing with a branch of a chain add –yl
d. Examples 1) 2)

37. 4. Rules for naming hydrocarbons
a. Name the longest chain (root)
1) Find the longest continuous chain
2) Select proper root
b. Name the compound type (suffix)
1) –ane ; -ene ; -yne
2) If the chain forms a ring, add ‘cyclo’

38. C. Name the branches
1) Each branch has a prefix and ends in –yl
2) Branches precede the chain name. If 2 or more branches, alphabetize the branches.
3) To specify where a branch occurs, number the carbon atoms starting at the end closer to a branch
4) If a double or triple bond is in the chain, the number of the carbon where the bond occurs precedes the root name