2.3 Atomic Theories

Greeks (5 th Century B.C.) – coined the term “atoms” to describe invisible particles of which substances were composed Aristotle (3 rd Century B.C.) – believed the universe was made of only 4 substances: Earth, Air, Water and Fire Earth, Air, Water and Fire John Dalton (1803) – Atomic Theory of Matter Matter composed of indivisible particles called atoms Matter composed of indivisible particles called atoms Elements contain identical atoms Elements contain identical atoms Different elements contain different atoms Different elements contain different atoms Atoms can combine from two or more elements to form new substances Atoms can combine from two or more elements to form new substances

J.J. Thompson (1897) – atoms contained negatively charged particles called electrons; envisioned a positive sphere with embedded electrons; sphere had a net charge of “zero”; termed the “Raisin Bun” Model H.Nagaoka (1904) – envisioned a positive sphere with a ring of electrons orbiting it (similar to the rings of Saturn) Ernest Rutherford (1914) – envisioned a very small positively charged nucleus surrounded by electrons; nucleus consisted of 1/1000 th of the total space of the atom

Niels Bohr (1921) – used Rutherford’s nuclear model with electrons ‘quantized’ in specific energy levels; became known as the Bohr- Rutherford Model (looked similar to planets orbiting the Sun) Erwin Schrodinger (1926) - Quantum Mechanics Theory – electrons were not in definite places, rather in “probability clouds”; similar to rotating fan blades James Chadwick (1932) – nucleus of the atom contained neutral particles called “neutrons”; had the same mass as protons and shared the nucleus with them

Isotopes Frederick Soddy (1913) – discovered the existence of isotopes Isotopes are a form of the same element in which the number of protons and electrons is the same, but the number of neutrons is different (example: carbon-12 and carbon-13) In other words, isotopes have the same atomic number but different atomic mass

atomic natural atomic natural #p #e #n mass abundance #p #e #n mass abundance Carbon % Carbon % Average atomic mass: a.m.u. Average atomic mass: a.m.u. Bromine % Bromine % Average atomic mass: a.m.u. Average atomic mass: a.m.u.

Bohr’s Theory of Atomic Structure Each electron in an atom have a fixed amount of energy related to the circular orbit in which it is found Electrons cannot exist between orbits, but they can move into unfilled orbits if a “quantum” of energy is absorbed or released The higher the energy level, the further it is from the nucleus The maximum number of electrons in the first three levels is: 2, 8, 8

Example: aluminum Atomic number: 13 (13 protons p +, 13 electrons e - ) Atomic number: 13 (13 protons p +, 13 electrons e - ) Electrons must be distributed amoungst 3 orbits around the nucleus using the 2,8,8 rule Electrons must be distributed amoungst 3 orbits around the nucleus using the 2,8,8 rule Diagram: Diagram: Al Al 3e- (3 rd level – “valence” level) 3e- (3 rd level – “valence” level) 8e- (2 nd level) 8e- (2 nd level) 2e-(1 st level) 2e-(1 st level) 13p+(nucleus) 13p+(nucleus)

Formation of Monatomic Ions Ions – atoms which have either gained or lost electrons to become stable; unlike atoms, ions always have a net charge The reason atoms gain or lose electrons to become ions is to attain a filled outermost (valence) shell Metals typically lose electrons to become positively charged (+); while non-metals typically gain electrons to become negatively charged (-) We will limit our discussion to the first 20 elements for simplicity reasons

Metal Ions Group 1 metals (e.g. Li, Na, K) donate one valence electron to become +1 ions donates to a non-metal to become… donates to a non-metal to become… E.g. sodium atom sodium ion 1e- (3 rd level) 1e- (3 rd level) 8e- (2 nd level) 8e- 8e- (2 nd level) 8e- 2e- (1 st level) 2e- 2e- (1 st level) 2e- 11p+ (nucleus) 11p+ 11p+ (nucleus) 11p+ Net Net Charge: 0 +1 Symbol: Na Na +

Metal Ions Group 2 metals (e.g. Be, Mg, Ca) donate two valence electrons to become +2 ions donates to a non-metal to become… donates to a non-metal to become… E.g. magnesium atom magnesium ion 2e- (3 rd level) 2e- (3 rd level) 8e- (2 nd level) 8e- 8e- (2 nd level) 8e- 2e- (1 st level) 2e- 2e- (1 st level) 2e- 12p+ (nucleus) 12p+ 12p+ (nucleus) 12p+ Net Net Charge: 0 +2 Symbol: Mg Mg 2+

Transition Metals Transition metals (groups 3-12) are very different from other metals in that their charges are much less predictable and often can have more than one ion charge (e.g. copper ions - Cu +, Cu 2+ )

Metal Ions Group 13 metals (e.g. Al) donate three valence electrons to become +3 ions donates to a non-metal to become… donates to a non-metal to become… E.g. aluminum atom aluminum ion 3e- (3 rd level) 3e- (3 rd level) 8e- (2 nd level) 8e- 8e- (2 nd level) 8e- 2e- (1 st level) 2e- 2e- (1 st level) 2e- 13p+ (nucleus) 13p+ 13p+ (nucleus) 13p+ Net Net Charge: 0 +3 Symbol: Al Al 3+

A note about Group 14 Since there are no metals in group 14 within the first 20 elements, we will move our discussion to non-metals Note: There are 3 metals in group 14 beyond the first 20 elements (Ge, Pb, Sn); however, their ion charges are somewhat unpredictable. We will treat them similar to the transition metals and look up their charges instead of trying to predict them

Non-metal Ions Group 15 non-metals (e.g. N, P) accept three valence electrons to become -3 ions accepts electrons from a metal to become … accepts electrons from a metal to become … E.g. phosphorus atom phosphide ion 5e- (3 rd level) 8e- 5e- (3 rd level) 8e- 8e- (2 nd level) 8e- 8e- (2 nd level) 8e- 2e- (1 st level) 2e- 2e- (1 st level) 2e- 15p+ (nucleus) 15p+ 15p+ (nucleus) 15p+ Net Net Charge: 0 -3 Symbol: P P 3-

Non-metal Ions Group 16 non-metals (e.g. O, S, Se) accept two valence electrons to become -2 ions accepts electrons from a metal to become … accepts electrons from a metal to become … E.g. sulfur atom sulfide ion 6e- (3 rd level) 8e- 6e- (3 rd level) 8e- 8e- (2 nd level) 8e- 8e- (2 nd level) 8e- 2e- (1 st level) 2e- 2e- (1 st level) 2e- 16p+ (nucleus) 16p+ 16p+ (nucleus) 16p+ Net Net Charge: 0 -2 Symbol: S S 2-

Non-metal Ions Group 17 non-metals (e.g. F, Cl, Br, I) accept one valence electron to become -1 ions accepts an electron from a metal to become … accepts an electron from a metal to become … E.g. chlorine atom chloride ion 7e- (3 rd level) 8e- 7e- (3 rd level) 8e- 8e- (2 nd level) 8e- 8e- (2 nd level) 8e- 2e- (1 st level) 2e- 2e- (1 st level) 2e- 17p+ (nucleus) 17p+ 17p+ (nucleus) 17p+ Net Net Charge: 0 -1 Symbol: Cl Cl -

Noble Gases Group 18 elements (e.g. He, Ne, Ar, Kr, Xe, Rn) were “born happy” will a filled outermost shell and therefore do not react with anyone E.g. argon atom 8e- (3 rd level) 8e- (3 rd level) 8e- (2 nd level) 8e- (2 nd level) 2e- (1 st level) 2e- (1 st level) 18p+ (nucleus) 18p+ (nucleus)

A note about hydrogen… Hydrogen is unique in that it can either GAIN or LOSE an electron to become stable donates to a non-metal to become… donates to a non-metal to become… E.g. hydrogen atom hydrogen ion e- (1 st level) e- (1 st level) p+ (nucleus) p+ p+ (nucleus) p+ Charge: 0 +1 Charge: 0 +1 Symbol: H H + Symbol: H H + accepts an electron from a non-metal to become… accepts an electron from a non-metal to become… E.g. hydrogen atom hydride ion e- (1 st level) 2e- e- (1 st level) 2e- p+ (nucleus) p+ p+ (nucleus) p+ Charge: 0 -1 Charge: 0 -1 Symbol: H H - Symbol: H H -

Homework Worksheet #4.

Ionic Compounds Ionic compounds are formed when metals donate electrons to non-metals Metals are left with a positive charge and are called cations (e.g. Na +, Mg 2+ ) Non-metals are left with a negative charge and are called anions (e.g. Cl -, N 3- )

Ionic Compounds Group 1 elements (Li, Na, K) react very readily with Group 17 elements (F, Cl, Br, I) because an exchange of one electron results in both ions having a filled outermost shell; e.g. sodium (atom) + chlorine (atom) sodium (ion) + chloride (ion) sodium (atom) + chlorine (atom) sodium (ion) + chloride (ion) NaCl Na + Cl - NaCl Na + Cl - e- 7e- 8e- e- 7e- 8e- 8e- 8e- 8e- 8e- 8e- 8e- 8e- 8e- 2e-2e-2e-2e- 2e-2e-2e-2e- 11p+17p+11p+17p+ 11p+17p+11p+17p+

Ionic Compounds Group 2 elements (Be, Mg, Ca) react very readily with Group 16 elements (F, Cl, Br, I) because an exchange of two electrons results in both ions having a filled outermost shell; e.g. calcium (atom) + oxygen (atom) calcium (ion) + oxide (ion) calcium (atom) + oxygen (atom) calcium (ion) + oxide (ion) Ca O Ca 2+ O 2- Ca O Ca 2+ O 2- 2e- 2e- 8e-8e- 8e-8e- 8e- 6e-8e- 8e- 8e- 6e-8e- 8e- 2e-2e-2e-2e- 2e-2e-2e-2e- 20p+8p+20p+8p+ 20p+8p+20p+8p+