chapter 2: atoms, ions and compounds problems: 2.1-2.80, 2.99-2.101, 2.104-2.105, 2.109-2.112
TRANSCRIPT
Chapter 2: Atoms, Ions and Compounds
Problems: 2.1-2.80, 2.99-2.101, 2.104-2.105, 2.109-2.112
Chapter 2 – Atoms, Ions and Compounds
2.1 The Rutherford Model of Atomic Structure (or Nuclear Model)• Joseph John (J. J.) Thomson (1897) carried out
experiments with cathode rays.– His research group determined the charge-
to-mass ratio of the particles in the rays.– They determined the particles were
composed of tiny, negatively charged subatomic particles electrons (e–)
Atomic Structure
• Eugen Goldstein (late 1880s)– Discovered canal (or anode) rays which
were composed of positively charged subatomic particles protons (p+)
– And decades later, James Chadwick won the Nobel Prize winner for his discovery (1935) neutron (n) = neutral subatomic particle
Plum-Pudding Model of the Atom
• Thomson proposed that the atom was a uniform sphere of positively charged matter in which electrons were embedded– electrons are like
raisins in a pudding of protons
The Nuclear Atom: Protons and the Nucleus
• Ernest Rutherford was a scientist who did many pioneering experiments in radioactivity
• He had members of his research group test Thomson’s Plum-Pudding Model using radioactive alpha (α) particles, basically helium atoms with a +2 charge and much bigger than an electron.
Rutherford's Alpha-Scattering Experiment
• Alpha (α) particles shot at a thin gold foil that’s only a few atoms thick– A circular detector is
set up around the foil to determine what happens to the α particles.
– If Plum-pudding Model was correct, the α particles should go through the foil like bullets through tissue paper.
Experimental results: Most of the a particles went straight through, but some were deflected or even bounced back!
Rutherford's Alpha-Scattering Experiment
Consider what the scientists expected to observe given each model.
Rutherford's Alpha-Scattering Experiment
Rutherford’s interpretation of the results• Most alpha (α) particles pass through foil
– Atom is mostly empty space with electrons moving around the space
• Some α are deflected or bounce back– Atom must also contain a dense region, and
particles hitting this region are deflected or bounce back towards source.
• Dense region = atomic nucleus (contains atom’s protons and neutrons)
That’s why this is called the NUCLEAR Model of the Atom
Rutherford's Alpha-Scattering Experiment
Rutherford also estimated the size of the atom and its nucleus:
An atom is 100,000 times (105 or 5 orders of magnitude) bigger than its nucleus.
nucleus (d~10-15 m)
atom (diameter ~10 -10 m)
Rutherford's Alpha-Scattering Experiment
Example 1An atom is 100,000 times (105 or 5 orders of magnitude) bigger than its nucleus. If a nucleus = size of a small marble (~1 cm in
diameter), indicate the length in meters then identify a common item that corresponds to that size for the following:
a. 10 times bigger = _____________ m = ________________________b. 100 times bigger = _____________ m = _______________________c. 1000 times bigger = _____________ m = ______________________d. 10,000 times bigger = ____________ m = _____________________e. 100,000 times bigger = ____________ m = ____________________
Properties of Protons, Neutrons, and Electrons
Subatomic Particle Charge Location Mass (amu)
Proton +1 inside nucleus 1.00728
Neutron 0 inside nucleus 1.00866
Electron -1outside nucleus 0.00055
2.2 Isotopes
• Each element always has the same number of protons, but the number of neutrons may vary. Atoms with different numbers of neutrons are called isotopes. – Carbon exists as carbon-12, carbon-13, and
carbon-14 where each carbon atom has 6 protons but 6, 7, or 8 neutrons respectively
• Isotopes are identified with an element name followed by the mass number– Examples: uranium-235 (U-235), carbon-12
(C-12), cobalt-60 (Co-60), etc.
Atomic Notation
• shorthand for keeping track of number of protons and neutrons in the nucleus– atomic number (Z): whole number of p+ =
number of e– in neutral atom– mass number (A): whole number sum of
protons and neutrons in an atom• Note: electrons contribute almost no mass to
an atom
EAZ
element symbol
mass number
atomic number
Example Problem
Isotope Mass number
# of protons
# of neutron
s
# of electron
s
strontium-86
92Mo
zinc-72
136Ba2+
31P3-
Complete the following table:
2.4 The Periodic Table of the Elements
Know which elements are metals, semimetals, nonmetals using the Periodic Table.
Properties of Metalloids (or Semimetals)• Have properties of metals and nonmetals• For example, silicon is shiny like a metal and
acts as a semiconductor.
Properties of Metals• conduct heat &
electricity• malleable: can be
flattened into thin sheets
• ductile: can be stretched into a wire
• Examples: aluminum, copper, gold
Properties of Non-Metals• dull appearance• brittle• non-conductor• Examples: carbon
(graphite in pencils), sulfur
SOLIDS, LIQUIDS, AND GASES: KNOW the physical state of each element at 25C!
At standard state conditions (25C and 1 atm):– Only mercury (Hg) and bromine (Br) are liquid– H, N, O, F, Cl, and all Noble gases (group VIIIA) are gases– All other elements are solids
2.5 Trends in Compound Formation
Chemical bond: holds atoms or ions together in a compound
COVALENT BOND and MOLECULES• Covalent bond: sharing of a pair of electrons by 2
nonmetal atoms• Two or more covalently bonded atoms form a
molecule• Molecule: basic unit of a compound of covalently
bonded atoms – Consider the HCl, H2O, NH3, and CH4 molecules
below– Note how the chemical formula gives the actual
number of each atom present in the following compounds:
Diatomic Molecules
Recognize these elements that exist as diatomic molecules (X2):
H2 N2 O2 F2 Cl2 Br2 I2
Consider these the “diatomic seven” since there are seven of them, and six of them form a 7 on the periodic table.
Ions• When atoms lose or gain electrons, they form
charged particles called ions.– Metals lose electrons positively charged
ions = cations– Nonmetals gain electrons negatively
charged ions = anions• Main-group elements generally form ions—i.e.
gain or lose electrons—to get the same number of electrons as a Noble gas.– Ions formed by main-group elements are
usually isoelectronic with—i.e., have the same number of electrons as—one of the noble gases!
Ions Formed by Main-Group Elements
Group IA elements +1 charge: Li+, Na+, K+, etc. (“+” = “+1”)
Group IIA elements +2 charge: Mg2+, Ca2+, Sr2+, Ba2+,etc.
Group IIIA elements +3 charge: Al3+
Group VA elements -3 charge: N3-, P3-
Group VIA elements -2 charge: O2-, S2-, Se2-
Group VIIA elements -1 charge: F–, Cl–, Br–, I–,etc.
Ionic Bonds and Ionic Compounds
Another type of chemical bond: IONIC BONDIons in an ionic compound are held
together by ionic bonds.• ionic bond: electrostatic attraction holding
together positively charged metal cations and negatively charged nonmetal anions
Ionic Bonds and Ionic Compounds
• Formula unit: most basic entity of an ionic compound (eg. NaCl, Al2O3, etc.)– The formula gives the ratio
of ions (not actual #).– The 3D representation of
NaCl at the right shows a network of Na+ (purple) and Cl– ions (green).
– The formula, NaCl, indicates a 1-to-1 ratio of Na+ ions and Cl– ions present, not the presence of only one ion of each.
• An ionic compound is a network of ions, with each cation surrounded by anions, and vice versa.
• To melt the solid, these bonds must be broken!• Ionic compounds have very high melting
points compared to molecules like water.
Ionic Bonds and Ionic Compounds
2.6 Naming Compounds and Writing Their Formulas• Every element can be identified using a name,
symbol, or atomic number.• Know the names and symbols for the first
18 elements on the Periodic Table as well as those elements included in Figure 2.17 on p. 57, and uranium (U). The Periodic Table that will be given on quizzes and exams will include only the element symbols, atomic number, and atomic mass. Spelling counts!
Ionic CompoundsCATIONS: positively charged ions
– Metal atoms lose valence electrons to form cations.
I. Groups IA, IIA, IIIA elements, silver, zinc and cadmium form only one type of ion each.
– Group IA elements +1 charge always (e.g. Li+=lithium ion)
– Group IIA elements +2 charge always (e.g. Mg2+=magnesium ion)
– Group IIIA elements +3 charge always (e.g. Al3+=aluminum ion)
– silver ion = Ag+; zinc ion = Zn2+; cadmium ion = Cd2+
II. The Stock system is used to name transition metals, Pb, and Sn forming different charges.
– iron (Fe), a transition metal, forms two different ions: Fe2+ and Fe3+
– lead (Pb), in Group IVA, forms two different ions: Pb2+ and Pb4+
Ionic Compounds
When a metal can form more than one ion, each ion is named:
element name (charge in Roman numerals) + ion
Pb = lead (II) ion Cu = copper (I) ion Fe = iron (II) ionPb = lead (IV) ion Cu = copper (II) ion Fe = iron (III) ion
+ 2+ 2+
2+ 3+ 4+
ANIONS: formed only by nonmetals
When a nonmetal forms an ion, it is named:
• element stem name + -ide suffix + ionO = oxygen atom O2– = oxide ionN = nitrogen atom N3– = nitride ion
Ionic Compounds
Polyatomic Ions
Know the formulas and names of the following polyatomic ions:
CrO42– = chromate ion
ClO– = hypochlorite ionC2H3O2
– = acetate ionNO3
– = nitrate ionClO2
– = chlorite ionPO4
3– = phosphate ionNO2
– = nitrite ionClO3
– = chlorate ionCN– = cyanide ionOH– = hydroxide ion
NH4+ = ammonium ion
CO32– = carbonate ion
Hg22+ = mercury (I) ion
Cr2O72– = dichromate ion
HCO3– = hydrogen carbonate
ionSO4
2– = sulfate ionMnO4
– = permanganate ionSO3
2– = sulfite ionClO4
– = perchlorate ionSCN– = thiocyanate ionO2
2– = peroxide ion
Naming Ionic Compounds
1. Get the individual ions for each compound2. CATION NAME + ANION NAME, minus “ion” Name of compoundZnS = __________________ ______________________________ individual ions
name of compound Fe2(CrO4)3 =___________________________BaSO4 =___________________________Ni(OH)3 =___________________________Hg(ClO4)2 =___________________________Cu3(PO4)2 =___________________________CdCO3 =___________________________PbO2 =___________________________ K2O2 =___________________________
Given the name of a compound, predict the formula:KNOW charges on ions formed by representative elements!KNOW how to use polyatomic ions and their charges when given to you!aluminum cyanide: __________________________________________
individual ions formula of compound
copper(I) sulfide:________________ ____________________________barium chlorate: ________________ ____________________________zinc phosphate: ________________ ____________________________cobalt(III) carbonate:________________ ____________________________silver nitrate:________________ ____________________________tin(II) permanganate:________________ ____________________________calcium dichromate:________________ ____________________________
Naming Ionic Compounds
Binary Molecular Compounds
# of atoms
Greek prefix
# of atoms
Greek prefix
1 mono 6 hexa2 di 7 hepta3 tri 8 octa4 tetra 9 nona5 penta 10 deca
Binary molecular compounds are composed of 2 nonmetals
NAMING: # of atoms of element indicated by Greek prefix before element name
1. For first element, Greek prefix + element name
2. For second element, Greek prefix + element name stem + "ide"
– If only one atom present, “mono-” is usually omitted, except for CO=carbon monoxide.
ExamplesExample: CO2 = carbon dioxide
CCl4 = _________________________
SF6= ____________________________
P4O10 =__________________________
Cl2O5= ____________________________
Some binary molecular compounds also have common names• e.g. everyone knows (or should know) H2O is
waterOther common compounds and names you must know:
NH3 = ammonia CH4 = methane
H2O2 = hydrogen peroxide
Acids
ACIDS: Aqueous solutions of a compound that releases H+ ions
– usually have H in front, physical state indicated as aqueous (aq)
– naming depends on the ion from which the acid forms
F– = fluoride ion HF(aq) = hydrofluoric acidNO2
– = nitrite ion HNO2(aq) = nitrous acidNO3
–= nitrate ion HNO3(aq) = nitric acid For some acids, the stem name changes: PO4
3– = phosphate ion H3PO4 (aq) = phosphoric acid
Example ProblemsCl–
CO32–
SO3
2–
C2H3O2
–
Br– SO4
2–
ClO–