Chapter 2 Atoms, Molecules and Ions

Problems from Chapter 2 Pages 52- 54 1 – 6, 9-12, 14, 16-19, 22, 24-34, 36-46, 48, 50, 60

EH Assignment Due Unit 3 Chemical Nomenclature Minimum score Sec 1 Names/Symbols of Elements 90 Sec 2 Molecules and Ions Sec 3 Cations and Anions 85 Sec 4 Ionic Compounds Sec 5 Acids, Bases, and Salts 80 Sec 6 Elements, Compound and Mixtures

EH Assignment Due Unit 9 Atomic Structure Minimum score Sec 1 Elementary Particles and Isotopes 90 Test on Chapters 1 and 2

Dalton’s Atomic Theory (1808) Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2.1

Dalton developed the atomic theory to explain The Law of Conservation of Mass The Law of Definite Proportions The Law of Multiple Proportions Laws are summaries of experimental observations. Theories are models devised to explain laws.

Laws of Mass Conservation & Definite Composition Law of Mass conservation: The total mass of substances does not change during a chemical reaction. Law of Definite ( or constant ) composition: No matter what its source, a particular chemical compound is composed of the same elements in the same proportions by mass.

Law of Conservation of Mass 16 X 8 Y + 8 X2Y Law of Conservation of Mass 2.1

LAW OF MULTIPLE PROPORTIONS If two elements can combine to form more than one compound, the masses of one element that combine with with a fixed mass of the other element are in the ratio of small whole numbers. Carbon monoxide 1.00 g C to 1.33 g O Carbon dioxide 1.00 g C to 2.66 g O

Law of Multiple Proportions 2 Law of Multiple Proportions 2.1

THE STRUCTURE OF THE ATOM Dalton thought that the atom was indivisible, but by the end of the 19th century there was evidence that the atom was composed of smaller particles. The electron was discovered in the 1890’s by J. J. Thomson using a Crooks tube (cathode ray tube). He was able to measure the ratio of the electric charge to mass of the electron to be -1.76 x 108 C/g. Later an American physicist, Millikan, measured the charge to be -1.60 x 10-19 C.

Fig. 2.4a

Fig. 2.4b

J.J. Thomson, measured mass/charge of e- (1906 Nobel Prize in Physics) 2.2

Thomson’s charge/mass of e- = -1.76 x 108 C/g Measured mass of e- (1923 Nobel Prize in Physics) e- charge = -1.60 x 10-19 C Thomson’s charge/mass of e- = -1.76 x 108 C/g e- mass = 9.10 x 10-28 g 2.2

The mass of the electron was much less than the mass of the lightest atom. This led to Thomson’s “Plum Pudding” model of the atom.

2.2

(Uranium compound) 2.2

Rutherford Experiment Bombarded a thin gold foil with high-energy alpha particles from radium. Most alpha particles went through unaffected. About 1 in 10,000 deflected through a large angle. Lead to nuclear model of atom.

(1908 Nobel Prize in Chemistry) particle velocity ~ 1.4 x 107 m/s (~5% speed of light) atoms positive charge and most of its mass are concentrated in the nucleus 2. light electrons are in the outer part of the atom. 2.2

Rutherford’s Model of the Atom If the atom is the Houston Astrodome atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m If the atom is the Houston Astrodome Then the nucleus is a marble on the 50 yard line 2.2

By the 1930’s it was known that the nucleus contained two subatomic particles - the proton and the neutron. The nucleus has a volume which is only a tiny fraction of the atom. The electrons occupy the outer part of the atom.

Subatomic Particles (Table 2.1) mass p = mass n = 1840 x mass e- 2.2

X H H (D) H (T) U Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei Mass Number X A Z Element Symbol Atomic Number H 1 H (D) 2 H (T) 3 U 235 92 238 2.3

2.3

Do You Understand Isotopes? How many protons, neutrons, and electrons are in C 14 6 ? 6 protons, 8 (14 - 6) neutrons, 6 electrons How many protons, neutrons, and electrons are in C 11 6 ? 6 protons, 5 (11 - 6) neutrons, 6 electrons 2.3

THE PERIODIC TABLE Elements with similar chemical and physical properties are grouped together in vertical columns called GROUPS or FAMILIES. Rows are called PERIODS. Group 1 (1A) - alkali metals Group 2 (2A) - alkaline earth metals Group 17 (7A) - halogens Group 18 (8A) - noble gases Transition elements (3-12) metals, nonmetals, metalloids

Alkali Earth Metal Noble Gas Halogen Alkali Metal Period Group 2.4

A diatomic molecule contains only two atoms A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2 H2O NH3 CH4 A diatomic molecule contains only two atoms H2, N2, O2, F2, Cl2, Br2, I2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 2.5

cation – ion with a positive charge An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. Na 11 protons 11 electrons Na+ 11 protons 10 electrons anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Cl- 17 protons 18 electrons Cl 17 protons 17 electrons 2.5

A monatomic ion contains only one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom OH-, CN-, NH4+, NO3- 2.5

How do we know how many electrons are lost or gained? For elements near either side of the periodic table the atoms tend to form ions with the same number of electrons as the nearest noble gas. What kind of ions do these elements form? I, Ba, Al, S, Rb, Fe

A Polyatomic Ion Fig. 2.22

How many protons and electrons are in Do You Understand Ions? How many protons and electrons are in Al 27 13 ? 3+ 13 protons, 10 (13 – 3) electrons How many protons and electrons are in Se 78 34 2- ? 34 protons, 36 (34 + 2) electrons 2.5

All metals form cations. 2.5

2.6

TA p49

TA p49

An empirical formula shows the simplest A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2O molecular empirical H2O C6H12O6 CH2O O3 O N2H4 NH2 2.6

The ionic compound NaCl ionic compounds consist of a combination of cations and an anions they only have empirical formulas, no molecular formula the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6

Formula of Ionic Compounds 2 x +3 = +6 3 x -2 = -6 Al2O3 Al3+ O2- 1 x +2 = +2 2 x -1 = -2 CaBr2 Ca2+ Br- 1 x +2 = +2 1 x -2 = -2 Na2CO3 Na+ CO32- 2.6

Some Polyatomic Ions (Table 2.3) 2.7

For our test you need to know: Ammonium Hydroxide Nitrate Sulfate Carbonate Bicarbonate Phosphate

This is an ionic compound. This is the ammonium cation. How can you tell if a compound is molecular or ionic?

Which of these are ionic and which are molecular? SiCl2, LiF, BaCl2, B2H6, KCl, C2H4, NH4NO3

Chemical Nomenclature Ionic Compounds (salts) often a metal + nonmetal (or polyatomic ion) anion (nonmetal), add “ide” to element name BaCl2 barium chloride K2O potassium oxide Mg(OH)2 magnesium hydroxide KNO3 potassium nitrate 2.7

Transition metal ionic compounds indicate charge on metal with Roman numerals FeCl2 iron(II) chloride 2 Cl- -2 so Fe is +2 FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide 2.7

TA p45

Name these: LiF, MgCl2, NaOH, Al2(SO4)3, NH4NO3, FeO Write formulas for: potassium carbonate, ammonium sulfide cobalt(II) bromide, aluminum phosphate

Molecular compounds (binary) nonmetals or nonmetals + metalloids common names H2O, NH3, CH4 element further left in periodic table is 1st element closest to bottom of group is 1st if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom last element ends in ide 2.7

Molecular Compounds HI hydrogen iodide NF3 nitrogen trifluoride SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride TOXIC! NO2 nitrogen dioxide N2O dinitrogen monoxide Laughing Gas 2.7

Practice exercises 2. 6 and 2 Practice exercises 2.6 and 2.7 Name NBr3 and Cl2O7 Write formulas for sulfur tetrafluoride dinitrogen pentoxide

2.7

An acid can be defined as a substance that yields hydrogen ions (H+) when dissolved in water. HCl Pure substance, hydrogen chloride Dissolved in water (H+ Cl-), hydrochloric acid An oxoacid is an acid that contains hydrogen, oxygen, and another element. HNO3 nitric acid H2CO3 carbonic acid H2SO4 sulfuric acid 2.7

2.7

If IO3- is the iodate ion name IO2-, HIO3, and HIO2

A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide 2.7