1. Atoms, Molecules, and Ions
Chapter 2
Atoms, Molecules, and Ions

2. Laws of Chemical Combination
Established from scientific observations
Used to establish other scientific theories
Include the following 3 laws:
Law of Conservation of Mass
Law of Definite Proportions
Law of Multiple Proportions

3. Law of Conservation of Mass
Lavoisier-mid 1700s
Total mass remains constant during a chemical reaction
Entire mass of reactants must be accounted for in the products
Also called a mass balance
11.1g H g O2 = 100.0g H20

4. Law of Definite Proportions
Joseph Proust - Late 1700s
All samples of a compound have the same composition
Same mass proportions of elements present
(Same atomic ratios of elements present)
Water
11.1% H 88.9% O
(2 atoms H, 1 atom O)
Also called law of constant composition

5. Law of Definite Proportions
Copper Carbonate, CH2Cu2O5
All samples of a compound have the same composition

6. Dalton’s Law of Multiple Proportions
John Dalton- Early 1800s
When two or more different compounds of the same two elements are compared, the masses of one element that combine with a fixed mass of the second element are in the ratio of small whole numbers.
Example:
CO: 12g Carbon reacts with 16g Oxygen
CO2: 12g Carbon reacts with 32g Oxygen
CO2 has 2X as much oxygen as CO
Ratio = 32:16= 2:1
2:1 is ratio of small whole numbers in Dalton’s Law

7. Atomic Theory of Matter
All matter is composed of atoms
All atoms of a given element are alike in mass, but
atoms of different elements differ in mass
Compounds are formed when atoms of different
elements unite in fixed proportions.
A chemical reaction involves atomic rearrangement
No atoms are created or destroyed

8. The Atom and Sub-Atomic Particles
Proton
large, positively charged particle in nucleus
Electron
small, negatively charged particle in orbit around the nucleus
Neutron
large, neutral particle in nucleus
Elements are not electrically charged
Must have equal numbers of protons and electrons
electron
proton
neutron
nucleus

9. Atomic Symbols Atomic Number (Z) Protons in a nucleus
Determines element identity
Located lower left on symbol
Mass Number (A)
# protons + # neutrons
Determines isotope identity
Located upper left on symbol

10. Isotopes
Elements with the same number of protons and electrons, but differing number of neutrons
Used in chemistry for structure identification or to follow a particular molecule through a reaction
Example: hydrogen and deuterium
Water, H2O
H has 1 proton and 0 neutrons in nucleus
Most abundant
Heavy water, D2O
D (deuterium) has 1 proton &1 neutron in nucleus
Occurs 1/6700 molecules 1 H 2

11. Atomic Mass Units (AMU)
Weighted average of the masses of the naturally
occurring isotopes of an element
Mass of protons & neutrons in nucleus
Mass proton or neutron: ~ 1.66 x g /particle
Particle mass independent of element
Electrons mass ignored
Atoms are difficult to weigh due to extremely small size
Use AMU- atomic mass units (u)
Internationally based on pure isotope C-12
C-12 atom contains 6 protons + 6 neutrons = 12 particles
1 u = weight of 1 particle = 1.66 x g
C-12 then has a mass of u

12. Atomic Mass Calculations
Percent abundance
isotope mass ÷ total mass
Fractional abundance
% abundance ÷ 100
Isotope contribution
fractional abundance x atomic mass
Atomic mass
sum of all mass contributions from isotopes

13. Calculating the Atomic Mass for C
Look up mass of isotopes
C u
C u
Look up % Abundances
C %
C %
Calculate contribution from Isotopes
C x u = u
C x u = u
Add up contributions to determine atomic mass
Contribution of C-12 + Contribution of C-13
11.87 u u = u

14. Mendeleev’s Periodic Table 1869
Arranged the known elements in order of increasing atomic weight from left to right and from top to bottom in groups.
Elements with similar properties are placed in the same column.
Used table to predict properties of some undiscovered Mendeleev elements

15. Germanium: Prediction vs. Observation

16. The Modern Periodic Table

17. Metals Characteristics Metallic, Conductive, Malleable,
Location: Left side of main table
Characteristics
Metallic, Conductive, Malleable,
Ductile, Luster, Solids (not Hg) ,
Positive ionic charge

18. Non-metals Location: Right side of table Characteristics:
Nonmetallic Brittle solids
Gas, liquids or solids Negative or uncharged ions

19. Noble Gases Location: Last column on the right side of table
Characteristics
Nonreactive Gases
Extremely stable Difficult to ionize

20. Halogens Location: Last column before noble gases Characteristics
Reactive 1- charge most common
Diatomic Stable ionized in water

21. Metalloids Location: Border between metals and nonmetals
Characteristics
Metal/nonmetal characteristics Solids
Variable ionic charges Semi-conductors

22. Lanthanides and Actinides
Table Location: Bottom 2 rows below table
Characteristics
Very reactive, Often radioactive, Unstable
Solids, Difficult to measure, and Positive ions

23. Molecules and Chemical Formulas
A group of 2 or more atoms held together with covalent bonds
Chemical Formula
Symbolic representation of molecular composition
Three different types of chemical formula
Empirical: ratio of atoms
Molecular: types and numbers of atoms
Structural: relationship of atoms in molecule

24. Chemical Formulas Acetic Acid: 2C 2O 4H
Empirical Formula
Shows ratio of atoms
CH2O
Molecular Formula
Shows number and type of atoms
C2H4O2
Structural Formula
Shows relationship between atoms
CH3COOH

25. Naming Binary Compounds
Name has 2 words, one for each element: N2O4
1st word is for 1st element: N: Nitrogen
2nd word is stem of element name O: Oxygen
change ending to “–ide” O: Oxide
Use prefixes to designate # of atoms
2N: Dinitrogen 4O: Tetroxide
Put it together N2O4: Dinitrogen tetroxide

26. Prefixes for Molecular Compounds Memorize!

27. Chemical Formulas for Binary Compounds
Choose the first element symbol Boron Trifluoride
Use the one farthest to left (metal) B
Lowest in group if in same group
Then write the other atom such as a nonmetal F
Remember the subscripts
Count # of atoms of each element 1B, 3F
BF3

28. Ionic Compounds (salts)
Atoms in a reaction may gain or lose electrons
Both atoms become charged and are called ions
Anion: atom gaining the electron is negatively charged
Cation: atom losing the electron is positively charged
The net charge is 0
Cations & anions balance out over the entire compound
No distinct molecular units
Positive charge of one ion attracts all nearby negative charges

29. Predicting Ionic Charge from the Periodic Table Metals, Metalloids and Nonmetals
Goal: Get to column 8A by going to the right or left
Right: Count each box as -1 until reaching 8A
Left: Count each box as +1 until reaching 8A in previous row
The correct charge is usually the smallest number
Left Side (metals): Li1+ or Li-7 Be2+ or Be6-
Right Side (nonmetals): O2-, F1-, He0, Ne0
Center (metalloids): B3+ (B5-), C4+ C4- , N5+, N3-

30. Predicting Ionic Charge from the Periodic Table Group B (Transition Metals)
Charge is positive, but unpredictable
Charge on a transition metal is designated with a
Roman numeral when naming the compound
Iron (III) oxide contains Fe3+

31. Naming Ions Add the word ION after element names
Sodium Chloride: NaCl Table Salt
Na Cl
loses 1 e- gains 1 e-
Na Cl-
Sodium ion Chloride ion
Net Charge = 0 = (+1) + (-1)

32. Monotomic and Polyatomic Ions
Atoms can lose or gain electrons singly or as a group
Monoatomic ions
Lose or gain electrons singly
The total charge is on a single atom
ex: Na+
Polyatomic ions
Lose or gain electrons as a group
The total charge is spread over 2 or more atoms
ex: SO4 2-

33. Polyatomic Ions Memorize the following polyatomic ions Ammonium NH4
+ Hydronium H3O+
Phosphate PO4
3- Acetate CH3COOHydroxide
OH- Nitrate NO3 -
Cyanide CN- Sulfate SO4 2-
Permanganate MnO4
- Chlorate ClO3
Carbonate CO3
2- Perchlorate ClO4

34. Acids, Bases and Salts Acid Characteristics taste sour sting the skin
turn litmus paper from blue to red
react with metals to produce ionic salts and hydrogen gas
Base Characteristics
taste bitter
feel slippery on the skin
turn litmus red to blue
react with acids to become neutral, often form ionic salts

35. Arrhenius Acids and Bases
Compound that ionizes in water to form a solution
of H+ ions and anions
Base
Compound that ionizes in water to form a solution of OH- ions and cations
Neutralization
Reaction between Arrhenius acid & base
H+ + OH- =H2O and cation + anion = salt
Uses
Identify acids and bases
Write formulas

36. Naming Arrhenius Acids
Binary Acids (HX)
Contain hydrogen and 1 other type of atom
Dissolved in water, so change hydrogen to hydro
Change “–ide” ending to “–ic acid”
example HCl: hydrochloric acid
Ternary and Oxoacids-
Binary acid that also contains oxygen
Oxoacids have oxygen as one of the atoms
Change “–ate” ending to “–ic acid”
example Nitric Acid: HNO3

37. Common acids to memorize
Hydrochloric Acid: HCl
Sulfuric Acid: H2SO4
Nitric Acid: HNO3
Carbonic Acid: H2CO3
Phosphoric Acid: H3PO4
Perchloric Acid: HClO4

38. Naming Arrhenius Bases
Named the same way as ionic compounds
All Arrhenius bases contain OH (hydroxide) in the
Name the metal first, then use word “hydroxide”
Memorize:
Sodium hydroxide: NaOH
Potassium hydroxide: KOH
Ammonium hydroxide: NH4OH

39. Polyatomic Ions with Oxygen
Suffixes
-ate
Most common form Chlorate ClO3-
-ite
1 less oxygen than “ate”
Chlorite ClO2-
Prefixes
Hypo-
1 less oxygen than ite ion Hypochlorite ClO per 1 more oxygen than -ate ion.
Perchlorate ClO4-

40. Hydrogen attached to ion
Use numerical prefixes for # of hydrogen atoms
mono, di, tri, etc.
Then add the word “hydrogen” before ion name
H2PO4-
Dihydrogen phosphate ion

41. Hydrates
A compound that is associated with a fixed number of water molecules is called a hydrate
Naming Hydrates
Use numerical prefixes
for # of water molecules
Then add the word “hydrate”
A dot shows H2O association
Water molecules are part of the mass
Copper(II)sulfate pentahydrate, CuSO4 ∙ 5H2O