1. Atoms, Molecules and Ions

2. Dalton’s Atomic Theory (1808)
All matter is composed of extremely small particles called atoms
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
Atoms cannot be subdivided, created, or destroyed (Law of conservation of mass)
Atoms of different elements combine in simple whole-number ratios to form chemical compounds (Law of constant composition)
In chemical reactions, atoms are combined, separated, or rearranged
John Dalton

3. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

4. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

5. Mass of the electron is 9.109 x 10-31 kg
1909 – Robert Millikan determines the mass of the electron.
Mass of the electron is x kg
The oil drop apparatus

6. Conclusions from the Study of the Electron
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

7. Rutherford’s Gold Foil Experiment
Alpha particles are helium nuclei
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

8. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
Conclusions
The nucleus is small
The nucleus is dense
The nucleus is positively charged
“Like howitzer shells bounding off of tissue paper!”

9. Atomic Particles Particle Charge Mass (kg) Location Electron -1
9.109 x (1/1836 amu)
Electron cloud
Proton +1
1.673 x (1 amu)
Nucleus
Neutron
1.675 x (1 amu)

10. The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons)
Electrons are found outside of the nucleus (the electron cloud)
Most of the volume of the atom is empty space

11. Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
Modern theory states:
Atoms of an element have a characteristic average mass which is unique to that element.

12. Modern Atomic Theory #2 Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

13. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.
Element
# of protons
Atomic # (Z)
Carbon 6
Phosphorous 15
Gold 79

14. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0
Nuclide p+ n0 e-
Mass #
Oxygen – 18 8 10 18
Arsenic – 75 33 75
Phosphorus – 31 15 16

15. Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen-1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

16. Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon =
Carbon =
Isotope
Symbol
Composition of the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%

17. Calculating Average Atomic Mass
Average atomic mass = S(% of each isotope)(mass of each isotope)
100
Example: If 69Ga and 71Ga occur in the percents 62.1 and 37.9, calculate the average atomic mass of gallium atoms.

18. Molecules
Two or more atoms of the same or different elements, ___________ bonded together.
Molecules are discrete structures, and their formulas represent each atom present in the molecule.
Benzene, C6H6

19. Examples of Molecules Examples of Molecules – Substance and Formula
Element or Compound
Description
Hydrogen (H2); Oxygen (O2); Nitrogen (N2); Fluorine (F2); Chlorine (Cl2); Iodine (I2); Bromine (Br2)
Elements
Diatomic
Water (H2O)
Compound
Polyatomic
Ammonia (NH3)

20. Covalent Network Substances
Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Groups of these substances are called __________.
Why Not??
Graphite Diamond

21. Ions
Atoms have equal numbers of protons and electrons so they have ________
When atoms gain or lose electrons, the proton:electron numbers are uneven causing a charged particle called an ____

22. Ions Cation: A ________ ion Anion: A _______ ion
Mg2+ (monatomic), NH4+ (polyatomic)
Anion: A _______ ion
Cl- (monatomic), SO42- (polyatomic)
Ionic Bonding: Force of attraction between oppositely charged ions.
Ionic compounds form crystals, so their formulas are written _________ (lowest whole number ratio of ions).

23. Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions
H+ Li+ Na+ K+
Group 2: Loses 2 electrons to form 2+ ions
Be2+ Mg2+ Ca2+ Sr2+ Ba2+

24. Predicting Ionic Charges
Group 13: Loses 3 electrons to form 3+ ions
B3+ Al3+ Ga3+
Group 14: Loses 4 electrons or gains 4 electrons
Caution! C22- and C4- are both called carbide

25. Predicting Ionic Charges
Group 15: Gains 3 electrons to form 3- ions
N3- P3- As3- Nitride Phosphide Arsenide
Group 16: Gains 2 electrons to form 2- ions
O2- S2- Se2- Oxide Sulfide Selenide

26. Predicting Ionic Charges
Group 17: Gains 1 electron to form 1- ions
F1- Cl1- Br I1- Fluoride Chloride Bromide Iodide
Group 18: Stable Noble gases do not form ions!

27. Predicting Ionic Charges
Groups 3-12: Many transition elements have more than one possible oxidation state (charge).
Iron(II) = Fe2+ Iron(III) = Fe3+
Some transition elements have only one possible oxidation state.
Zinc = Zn2+ Silver = Ag+

28. Binary Compounds of Nonmetals and Metals (Ionic Compounds)
To write the formula from the name
Write the formulas for the _____ and _____, including CHARGES!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using ________. Use parentheses if you need more than one of a polyatomic ion.

29. Writing Ionic Compound Formulas
Example: Barium nitrate
Example: Ammonium sulfate
Example: Iron(III) chloride
Example: Aluminum sulfide
Example: Magnesium carbonate
Example: Zinc hydroxide
Example: Aluminum phosphate

30. Naming Ionic Compounds
1. Cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride

31. Naming Ionic Compounds
Metals with multiple oxidation states
Some metal forms more than one cation
Use Roman numeral in name
PbCl2
Pb2+ is the lead(II) cation
PbCl2 = lead(II) chloride

32. Binary Acids
An acid is a compound that produces hydrogen ions in water
Formed when hydrogen ions combine with monatomic anions
To name an acid
Use the prefix hydro followed by the non-metal name modified to an –ic ending
Add the word acid
Example: HCl and HI

33. Polyatomic Ions and Oxyanions
Polyatomic ion – more than one element are combined to create a species with a charge
Have to memorize these
Oxyanions – a polyatomic ion that contains oxygen
To name a compound containing a polyatomic ion
Positive ion written first
Polyatomic ion written second, unmodified
Example: K2CO3

34. Oxyacids Formed when hydrogen ions combine with polyatomic oxyanions
To name an oxyacid
Use the name of the oxyanion and replace the –ite ending with –ous or the –ate ending with –ic
Add the word acid
Example: HClO4 and HNO3

35. Naming Binary Compounds of Two Nonmetals
Two nonmetals form molecules
First element in the formula is named first.
Second element is named as if it were an anion.
Use prefixes – mono, di, tri, tetra, penta, hexa, hepta and octa
Only use mono on second element -
P2O5 = diphosphorus pentoxide
CO2 = carbon dioxide
CO =carbon monoxide
N2O = dinitrogen monoxide

36. Hydrates Ionic formula units with water associated with them
Water molecules are incorporated into the solid structure of the ions
To name a hydrate
Use the naming of an ionic compound followed by the term hydrate with an appropriate prefix for the hydrate
Example: CuSO4.5H2O