Atoms, Molecules and Ions

1
Chapter 2

• Atoms, Molecules and Ions

2
Objectives

• Recall Daltons Atomic Theory.
• Explain the discovery of electrons and the
  nucleus.
• Use atomic mass and atomic number in
  calculations.
• Discuss periodic table setup.

3
Fundamental Chemical Laws

• Conservation of Mass Lavoisier
• Mass is neither created or destroyed
• Mass of reactants equals mass of products
• Definite Proportions Proust
• A given compound always contains the same
  proportions of elements by mass
• Formulas of compounds dont change

4
Cont.

• Multiple Proportions Dalton
• When two elements form a series of compounds, the
  ratios of the masses of the second element that
  combine with one gram of the first element always
  can be reduced to small whole numbers
• Two elements can form different compounds
• CO and CO2

5
Daltons Atomic Theory - 1808

• Each element is made up of tiny particles called
  atoms
• The atoms of a given element are identical, the
  atoms of different elements are different
• Chemical compounds are formed when atoms combine
  with each other. A given compound always has the
  same relative numbers of and types of atoms

6
Cont.

• Chemical reactions involve reorganizations of
  atoms. The atoms themselves are not changed in a
  chemical reaction
• 1 and 2 are not quite true
• Subatomic particles
• Isotopes
• Also made first table of atomic masses

7
Discovery of the Electron

• JJ Thomson 1898 to 1903
• Study of cathode rays
• Negatively charged b/c they are attracted to a
  positive field
• All metals create the same type of rays
• Determined the charge to mass ratio
• -1.76x108 C/g

8
Cathode Ray Tube
9
Thomsons Conclusions

• All atoms contain electrons
• Negatively charged
• Reasoned atoms must contain positive charge
• Proposed the plum pudding model of the atom
• Positive goo with negative charges embedded

10
Plum Pudding Model
11
Robert Millikan

• Used oil drop experiment to determine the mass of
  an electron (9.11x10-31 kg)

12
Radioactivity

• Alpha Particles (a)
• Nuclear Particle with a 2 charge
• Helium nucleus
• Little danger to life
• Beta Particles (ß)
• High speed electrons
• Gamma Particles (?)
• High energy light
• Most dangerous to life

13
Discovery of Nucleus

• Ernest Rutherford 1911
• Used Gold Foil Experiment
• Shot Alpha Particles at gold foil surrounded by
  screen to detect alpha particles
• He expected most particles to go through the foil
  
• Most when straight through
• Some deflected

14
Gold Foil Experiment
15
Conclusions

• As though a Howitzer bounced off a tissue
• Atom has something in it that is
• Very dense
• Positively charged
• Proposed nuclear model of the atom

16
Plum Pudding vs. Nuclear Model
17
Modern Atom

• Consists of Nucleus and Electron Cloud
• Nucleus - Protons and Neutrons
• Small, Dense, and Positively Charged
• Diameter of 10-13m
• 250 million tons/ cubic centimeter!
• Electron Cloud - Electrons
• Much larger than the nucleus
• Diameter of 10-8m
• 5 Powers of 10 different!

18
Mass and Charge
Particle Mass Charge
Proton 1.673x10-27 kg 1
Neutron 1.675x10-27 kg 0
Electron 9.109x10-31kg -1
19
Counting Atomic Parts

• Atomic Number Number of protons
• Determines the type of element
• Same as of electrons in neutral atoms
• Atomic number of Nitrogen
• Atomic number of Potassium
• Mass Number Number of protons neutrons
• Both numbers are always whole numbers!

7
19
20
Isotopes

• Same number of protons different number of
  neutrons
• Isotopes are distinguished by name or symbols
• Hyphen Notation
• Carbon -12 or Carbon - 14
• Symbols

Mass Number
Atomic Number
21
Example

• Determine the number of protons, neutrons, and
  electrons in Sulfur-32
• Protons 16
• Neutrons 16
• Electrons 16

22
Example

• Determine the number of protons, neutrons, and
  electrons in
• Protons 33
• Neutrons 42
• Electrons 36

23
Example

• Write the symbol for a neutral atom with 21
  electrons 23 neutrons

24
Periodic Table Basics

• Rows are called Periods
• Columns are called Groups or Families
• Group 1 Alkali Metals (Lose 1 Electron, 1)
• Group 2 Alkali Earth Metals (Lose 2 Electrons,
  2)
• Group 17 Halogens (Gain 1 Electron, -1)
• Group 18 Noble Gases (No charge)
• Metals are on the left side of the table
• Nonmetals are on the right

25
(No Transcript)
26
Homework

• p. 75 s 46,50,58

27
Objectives

• Provide names for ionic, covalent, and acidic
  compounds
• Provide formulas for ionic, covalent, and acidic
  compounds

28
Ionic Compounds

• Always involves a transfer of electrons
• Atoms gain or lose electrons to achieve the same
  number of electrons as noble gases
• Metals lose electrons
• Become positive ions (cations)
• Nonmetals gain electrons
• Become negative ions (anions)

29
(No Transcript)
30
Naming Ionic Compounds

• Cation is named first, anion second
• Cation keeps its name, anion ends in ide
• Ex. Sulfide, Chloride, Oxide
• CaCl2
• Calcium Chloride
• NaBr
• Sodium Bromide
• Ba3P2
• Barium Phosphide

31
Naming Ionic Compounds II

• Some cations can have more than one charge
• Transition metals, Tin, Lead
• A roman number is placed after the cation name to
  indicate charge
• CuI2
• Copper (II) Iodide
• CuI
• Copper (I) Iodide

32
Cont.

• Exceptions to the transition metal rule
• Silver, Zinc, Chromium only form one ion
• Silver only 1
• Zinc only 2
• Chromium only 2
• Therefore roman numeral is not required
• AgBr
• Silver Bromide not Silver (I) Bromide

33
Polyatomic Ions

• A covalently bonded group of atoms that possess a
  charge.
• Table 2.5 MUST be memorized!
• One more oxygen that normal Per- -ate
• Normal number of oxygens -ate
• One less oxygen than normal -ite
• Two less oxygens than normal Hypo- -ite

34
(No Transcript)
35
Naming Ionic Compounds III

• Name cation as normal
• Include roman number if necessary
• Name polyatomic ion
• NaNO3
• Sodium Nitrate
• NiSO4
• Nickel (II) Sulfate

36
Naming Acids

• Two types of acids
• Hydrogen and halogen (Hyrdohalic Acids)
• Named Hydroblaic Acid
• HCl Hydrochloric Acid
• Hydrogen and polyatomic ions
• Name depends on name of polyatomic ion
• -ate becomes ic
• -ite becomes ous

37
Cont.

• H2SO4
• Sulfuric Acid
• HNO3
• Nitric Acid
• H2SO3
• Sulfurous Acid
• HNO2
• Nitrous Acid

38
Covalent Compounds

• Use Number prefixes
• 1 Mono 6 Hexa
• 2 Di 7 Hepta
• 3 Tri 8 Octa
• 4 Tetra 9 Nona
• 5 Penta 10 Deca

39
Cont.

• The first element gets its name
• If something other than one use a prefix
• The second element ends in ide
• Use prefix no matter what
• CO
• Carbon Monoxide
• CO2
• Carbon Dioxide

40
Writing Formulas from Names

• Work backwards from the name
• Determine charges of the ions
• Make ratio so charges balance
• Reduce formulas to lowest ratio
• Except formulas that are fixed
• Peroxide O2-2

41
Examples

• Calcium Fluoride
• CaF2
• Calcium Nitrate
• Ca(NO3)2
• Iron (III) Sulfate
• Fe2(SO4)3
• Titanium (IV) Oxide
• TiO2

42
Homework

• p. 77 's 60,62,64,68 a-f,70 a-f