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Unit 2 Atoms, Molecules, & Ions.

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1 Unit 2 Atoms, Molecules, & Ions

2 Law of Constant Composition (Definite Proportion): All samples of a compound have the same composition, the same proportion by mass of the constituent elements. H2O by mass is always % H 88.81 % O A 0.1 g sample of Mg, when combined with O2 (g) yields g MgO. A second sample, g of Mg is also combined with O2 (g). What is the mass of MgO produced?

3 Law of Multiple Proportions: When 2 elements form a series of compounds, the ratio of the masses of the second element that combine with 1.0 g of the first element can always be reduced to small whole numbers. The following data were collected for several compounds of nitrogen combining with 1 g oxygen. Determine the formulas. Compound A = g N Compound B = g N Compound C = g N

4 Dalton’s Atomic Theory
Each element is made up of tiny particles called atoms. The atoms of a given element are identical; atoms of different elements are different in some fundamental way. Chemical compounds are formed when atoms of different elements combine with each other. They combine in simple whole number ratio, always with the same relative number and types of atoms. Chemical reactions involve reorganization of the atoms, changes in the way they are bonded together. The atoms themselves are indestructible.

5 History Need to know theory, NOT dates or names! 1803 John Dalton discovered that elements are made of atoms. He thought that atoms were solid, like a marble. Early 1800’s Michael Faraday discovered that a cathode ray is negatively charged particles 1874 George Stoney coined the term “electron” 1875 Crooks discovered the electron has a mass of 9.1 x g and have a negative (-) charge. 1897 JJ Thomson used a cathode ray tube to discover negative ray consisting of electrons with a charge to mass ratio c/m = x 108 c/g 1895 Roentgen discovered when using a cathode ray tube, that there was fluorescence in the room. He found X rays. 1909 Robert Millikan worked with an oil drop to determine the magnitude of e- charge to be – x coulomb.

6 discovered radioactivity a alpha particle is a helium nucleus 4He
1896 Henri Bacquerel (1852 – 1908) discovered radioactivity a alpha particle is a helium nucleus 4He b beta particle is an electron e- g gamma ray is pure energy 1911 Earnest Rutherford (1871 – 1937) discovered that there is a mass in the center of the atom with a positive charge containing p+, the e- are around the nucleus, with the atom mostly empty space. 1932 Chadwick discovered the neutron which is also found in the nucleus. 2

7 Electrical charge (coulomb)
Subatomic particles Particle amu mass (g) Atomic charge Electrical charge (coulomb) Proton (p+) 1.0073 1.673 x 10-24 +1 x 10-19 Electron (e-) 9.109 x 10-28 -1 x 10-19 Neutron (n) 1.0087 1.675 x 10-24

8 Nucleus has a density of 1 x 1013 to 1 x 1014 g/cm3
There are 4 forces that hold an atom together: 1. gravitational – attraction between 2 bodies with mass 2. electromagnetic – attractive and repulsive forces due to charged or magnetic objects 3. strong nuclear force – keep protons which all have positive charges from flying apart 4. weak nuclear force – responsible for (b) beta decay in radioactivity

9 C Periodic Table 6 12.0111 Atomic mass Element symbol Atomic number
Elements with the atomic mass in ( ) means that very little of it exists at any given time. Scientists have not been able to attain an good average atomic mass.

10 Nuclear Notation A X Atomic mass = p+ + N Atomic number = p+
Isotopes = Atoms of the same element (so the same atomic number) but with a different mass because they have a different number of neutrons. Chemical symbol Z Atomic number

11 Alpha decay a A X  A-4 N + 4 He
Kinds of Decay Alpha decay a A X  A-4 N + 4 He N = new element Beta decay b A X  A N + e- N  p+ + e- Z Z-2 Z Z+1

12 Atomic Mass is the average of the number of protons and Neutrons found in the nucleus
Average Atomic Mass = S ( ) fraction of abundance of each isotope ( ) Mass of each isotope In naturally occurring Argon, % of the atoms are 40Ar with a mass of amu (u), 0.337% 36Ar with a mass of u, and 0.063% 38Ar with a mass of u. Calculate the Average mass of Argon. 18 18 18

13 Classification of elements as metals, nonmetals, and metalloids.

14 Alkali Earth metals +2 ion
Noble gases Halogens -1 ion Alkali metals +1 ion Alkali Earth metals +2 ion Chalcogens -2 ion Transition metals Rare Earth Metals

15 Periodic table with atomic symbols, atomic numbers, and partial electron configurations.

16 Diatomic molecules There are 7 elements that do not exist as 1 atom alone. As an element, they are always as 2 atoms bonded together as a molecule. Location of the diatomic elements

17 Molecular Compound – molecule with more than 1 type of atom
Molecular Formula – shows the actual # of atoms Empirical Formula – shows the relative # of atoms in the smallest possible whole # ratio Structural Formula – shows how the atoms are assembled into the molecule

18 Chemical Bonds – the forces holding atoms together
Ionic bonds – transfer of e- resulting in ions attracted to each other Ions – atom or group of atoms which have gained or lost e- (s), therefore carrying an electrical charge. cation (+) positive charged ion anion (-) negative charged ion Covalent bonds – sharing of e- Polar covalent bond – an uneven sharing of e- resulting in a molecule which appears to have opposites charges on different parts of the molecule. Nonpolar covalent bond – an even sharing of e-

19 Naming compounds Inorganic binary compounds type I
one atom (+) cations maintain atom name one atom (-) anion changes name (-ide) (+) cation is followed by the (-) anion Inorganic type II – some anions form multiple charges Use roman numerals to indicate charge example: CuCl – copper (I) chloride, CuCl2 – copper (II) chloride Or/ higher charge name ends with (-ic) while lower charge name ends with (-ous) example: FeCl3 ferric chloride FeCl2 ferrous chloride

20 Element Latin naming Copper Cu+2 Cuprous, Cu+3 Cupric Iron Fe+2 Ferrous, Fe+3 Ferric Mercury Hg2+2 Mercurous, Hg+2 Mercuric Lead Pb+2 Plumbous, Pb+4 Plumbic Tin Sn+2 Stannous, Sn+4 Stannic Chromium Cr+2 Chromous, Cr+3 Chromic Manganese Mn+2 Manganous, Mn+3 Manganic Cobalt Co+2 Cobaltous, Co+3 Cobaltic

21 Polyatomic – must memorize the names!
Oxy anions: elements with oxygen acting as polyatomic Less than small number of oxygen (hypo-) Small number of oxygen (-ite) Large number of oxygen (-ate) More than large number of oxygen (per-) Example: hypochlorite ClO-1 chlorite ClO2-1 chlorate ClO3-1 perchlorate ClO4-1

22 Naming Acids – when dissolved in H2O they produce free protons (H+)
If anion does not contain oxygen Use prefix (hydro-) with suffix (-ic) example HS hydrosulfuric acid If anion does contain oxygen If anion ends in (-ate) use root name with (-ic) H2SO4 sulfuric acid If anion ends in (-ite) use root name with (-ous) H2SO3 sulfurous acid When multiple oxyanion use Example HClO4 perchlorate perchloric acid HClO3 chlorate chloric acid HClO2 chlorite chlorous acid HClO hypochlorite hypochlorous acid

23 Binary Covalent type III – 2 nonmetals
1st element in formula named first with full element name 2nd element named as if anion Prefixes used to denote number of atoms present Mono is never used for naming 1st element mono hexa di hepta - 7 tri octa tetra nona penta dec

24 Naming Organic Molecules
Name root name according to the number of carbons in the longest continuous chain. Name the suffix according to the bonding, carbon to carbon. Single bond C – C end with -ane Double bond C = C end with -ene Triple bond C C end with -yne Name substitutions on the main chain before the root name, giving the lowest possible number.

25 Hydrocarbon

26 Substitute Naming Group Names Methyl -CH3 Ethyl -CH2CH3
Propyl -CH2CH2CH3 Butyl -CH2CH2CH2CH3 Halides Chloro -Cl Iodo -I Bromo -Br Fluoro -F Prefixes used when more than one group of the same kind is attached di two substitutes tri three tetra- four penta- five hexa six

27 Naming organic molecules containing functional groups
Drop the e, add Functional Group Group Name Name ending - O H Alcohol ol Carboxylic Acid oic Acid Ketone one

28 Functional Group Group Name Name ending. Aldehyde - al
Functional Group Group Name Name ending Aldehyde al Double bond ene Triple bond yne Drop –ane, add

29 Small molecules only have one arrangement possible.
Isomers: Molecules with the same chemical formula, but different arrangement of atoms. With different arrangement of atoms, the properties of the chemicals are different. Small molecules only have one arrangement possible. CH4 Methane C2H6 Ethane

30 C5H12 has 3 isomers

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