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Study Guide – Fall Final
Physical Science – Dr. Hazlett
NOTE: THIS IS A STUDY GUIDE. IT MAY NOT CONTAIN ALL INFORMATION COVERED IN CLASS OR THAT MAY APPEAR ON THE EXAM. USE THE WORKBOOK INTRO PAGES TO EACH CHAPTER COVERED, THE POWERPOINTS, WORKSHEETS, THE TEXTBOOK, AND MOST IMPORTANTLY, YOUR NOTES TO STUDY!!!!!
I. Measuring Matter
A. Inertia – ability of an object to resist change/motion
B. Mass – the amount of matter in an object. NOT the same as weight
1. Mass (m) is a measure of inertia
2. Conservation of Mass – matter must be and is conserved, it can not be created nor destroyed, only transformed
3. Law of Definite Proportions – in given cmpd, same elements present in same proportions by mass
4. Conservation of Energy – Energy can not be created nor destroyed, only transformed
C. Volume (V) – the amount of space an object takes up, the cubic measure
1. Note: 1 cm3 = 1 mL; 1L = 1000 cm3; 1000 mL = 1 L
D. Density – the amount of matter in a given volume
1. Density (D) = Mass
Volume
2. D related to an object’s buoyancy
E. Weight – the measure of gravitational force (g) upon a mass
1. Weight can change based upon a Δ in g
2. wt. = mass x gravitational force
a. On earth > g = 9.81 m/s2
F. Mole – a constant for measurement
1. Avogadro’s Number (NA)
a. NA = 6.022314 x 1023 particles
II. Matter Properties
A. Types of Matter
1. Substance –
a. Element – single, pure substance (on Per. Table)
b. Compound – combination of elements, may be molecules
2. Mixture –
a. Heterogeneous – separable mixture
b. Homogeneous – nonseparable mixture
c. Separation Methods
(1) Filtration
(2) Distillation
(3) Crystallization
d. Miscible – mixes, immiscible – does not mix
3. Solutions – Solute + Solvent
a. Solute – What is Dissolved
b. Solvent – What the Solute is Dissolved In
c. Electrolyte – conducts electricity
B. Properties of Matter
1. Physical –
a. Includes density, color, temperature, boiling/melting points, and related
factors
2. Chemical – reactions (rxn)
a. Change in make-up of atoms or compounds, energy use and/or release
III. Temperature and Pressure
A. Temperature is always given in oC or oK
B. Pressure (P)
1. Measured in Atmospheres (atm) = 760 mmHg = 760 Torr = 101 325
Pascals (Pa or in N/m2) = 10.1325 kPa (or in N/cm2)
IV. States of Matter
A. Intermolecular Force (IMF) –
1. Kinetic Theory of Matter
a. All particles in matter are moving to some extent
b. Effect of Temperature and Pressure on Matter and its Particles
B. States (know characteristics of each and connected concepts)
1. Solid (s)
a. Crystalline – organized
b. Amorphous - random
2. Liquids
a. Viscosity – measure of flow
b. Surface Tension
c. Capillary Action – flowing uphill (paper towel example)
d. Vapor Pressure – evaporation over liquid
3. Gas
a. Pressure – force of particles colliding with walls of container
b. Diffusion – dispersion of the particles throughout the container
4. Plasma – atoms lose e-, highest KE
5. Bose-Einstein Condensate (BEC) – at absolute zero!, lowest level KE
6. Dark Matter
V. Changes in States of Matter
A. Phases
1. Phase Diagrams (P and T are Independent Variables)
a. Triple Point
b. Critical Point
B. Phase Change Terms
1. Solid to Liquid =melting/liquification
2. Liquid to Solid=freezing/solidification
3. Liquid to Vapor=evaporation
4. Vapor to Liquid =condensation
5. Solid to Vapor=sublimation
6. Vapor to Solid=deposition
The Atom and Periodic Table
I. The Atom
A. Nucleus
1. p+ gluon (strong force) no
2. Each p+ and n0made of 3 quarks held together by bosons (weak force)
3. Nucleus held to Lepton / e- by EMF and gauge bosons
4. Neutron decay causes beta radiation (e- and neutrinos)
a. Alpha Radiation – He nucleus (2 p+ and 2 n0)
b. Gamma – photons and other subatomic particles
B. Ions
1. p+ > e-; a positive ion called cation
2. p+ < e-; a negative ion called anion
C. Electron Models
1. Thomson and Plum Pudding Model
2. Rutherford and Planetary Model
3. Bohr and Quantum Model
a. 7 energy levels for e- orbits
D. Schrodinger and Electron Cloud Model
1. Heisenberg Uncertainty Principle
2. Electrons travel at speed of light
II. The Electron
A. Energy levels (n)
1. 7 Energy levels (Bohr Model)
2. e- in their lowest level called the ground state
3. When absorb energy, move up a level to their excited state
4. When release this energy, it leaves in the form of a photon.
a. This is a packet of light with a certain color (wavelength)
B. Sublevels or Orbitals
1. Each energy level (n) has a certain number of sublevels
2. 4 basic shapes of these sublevels
a. s = sharp = spherical
b. p = principle = lazy eight shape
c. d = diffuse = 4 leaf clover shaped
d. f = fundamental = six leaf clover shaped
3. These in a 3-dimensional, intertwined way make the e- cloud
4. Know valence energy and sublevels by use of periodic table
C. Aufbau Diagram
1. Aufbau Principle - fill lowest energy levels first
2. Pauli Principle – each suborbital (box) has maximum of 2 e- with
opposite spins
3. Hund’s Rule (Share the Cookies Rule) – each box must get one e-
before any one gets a second e-
III. Valence Orbitals / Blocks
A. Valence e- are the outer most shell of electrons
1. Octet Rule – atoms want to have a complete/full valence shell
a. Groups I and II - sb. B groups/Transition Metals - d
c. Lanthanides and Actinides - fd. Groups III through VIII - p
IV Organizing the Elements – History
A. Mendeleev – organized table by atomic weight
B. Moseley – organized table by atomic number
V. Table
A. Periods are the rows and represent the energy level (n) of valence e-
B. Groups/Families are columns and have same # valence e- for A elements
1. 10 elements have symbols that don’t match the first letters of name - know
C. Metals in A, and those in B are Transition Metals
1. Metals are to left of stair line (includes Al)
2. Lanthanides and Actinides are Transition Metals (Period 6 and 7)
D. Nonmetals – to right of stair line (Halogens and Noble Gases)
E. Metalloids – have side along stair line (Excludes Al and most times, B)
Group I A. Alkali Metals (Group 1)
- form ions, each with a +1 charge
II A. Alkaline Earth Metals (2)
- tend to lose two electrons per atom, forming ions with a +2 charge
The B Groups. Transition Metals (3-12)
- consist of metals in groups 3 through 12
- contain one or two valence electrons, in d or f blocks
- tend to have two or more common + oxidation states
- may form complex ions
III A. Boron Family (13)
- do not occur elementally in nature
- have three valence electrons in p block
- form +3 ions
- are metallic (except boron, which is a solid metalloid)
IV A. Carbon Family (14)
- includes a nonmetal (carbon), 2 metalloids (silicon and germanium) and 2 metals (Sn and Pb)
- occur in nature in both combined and elemental forms
- have four valence electrons in p block
V A. Nitrogen Family (15)
- consists of two nonmetals (nitrogen and phosphorus), two metalloids (arsenic and antimony), and one metal (bismuth)
- have five valence electrons in p block
- tend to form covalent compounds
- most commonly with oxidation numbers of +3 or +5
VI A. Oxygen Family (16)
- consists of three nonmetals (oxygen, sulfur, and selenium), one metalloid (tellurium), and one metal (polonium)
- have six valence electrons in p block
- tend to form covalent compounds with other elements
- tend to exist as diatomic and polyatomic molecules, such as O2, O3, S6, S8, and Se8
- commonly exist in compounds with the -2 oxidation state
VII A. Halogen Family (17)
- are nonmetals and occur in combined form in nature, mainly as metal halides
- form salts when react with alkalines
- have seven valence electrons, forming -1 ions
VIII A. Noble Gases (18)
- not reactive
- have a full outer energy level, completed octet rule (p block)
- are all gases
- are all nonmetals
Hydrogen
- metal and nonmetal, not in any group
- proterium, deuterium and triterium isotopes
IV.Metallic Character
1. The stair-step line divides the periodic table into metals, nonmetals, and
metalloids.
2. Metals lie below and to the left of the stair-step line. They include the elements
aluminum and polonium that border the stair-step line.
3. Properties of metals include:
a. Shiny luster
b. Conductivity of heat and electricity
c. Malleability (can be rolled into thin sheets)
d. Ductility (can be pulled into thin wires)
e. High melting point
f. Low first ionization energy
g. Form ionic compounds with nonmetals
h. Form basic oxides
Elements, Periodic Table and Electrons
I. The Formation of the Elements
A. Nucleosynthesis
1. H isotope – deuterium
a. Extreme temperatures and pressure/density
2. H He C fusion into final element Fe
B. Star Life Cycle
1. Nebula
2. Sun/Star
a. Balance between fusion and gravity
3. Red Giant
4. White Dwarf
5. Supernova
6. Neutron Star and/or Black Hole
II. Organizing the Elements – History of the Table
A. Antoine Lavoisier
B. Johann Dobereiner and the Law of Triads
C. John Newlands and the Law of Octaves
1. Both ideas fall apart due to exceptions/new elements discovered
D. Kekule and Meyer
1. Valency – outer electrons and orbits
D. Mendeleev – organized by atomic weight and properties, predictions
E. Moseley – organized by atomic number found via X-rays
F. Seaborg
1. Transuranium elements
III. The Periodic Table
A. Periods are the rows and represent the energy level (n) of valence e-
Valence Energy Level (n)Valence ShapesMax. Val. e-
1s 2
2s, p 2, 6 = 8
3s, p, d 2, 6, 10 = 18
4 s, p, d, f 2, 6, 10, 14 = 32
5 s, p, d, f 2, 6, 10, 14 = 32
6s, p, d, f 2, 6, 10, 14 = 32
7s, p, d, f 2, 6, 10, 14 = 32
- Shapes
- S is sphere
- P is principle or a figure 8
- D is diffuse – four leaf clove
- F is fundamental or 3d six leaf clover
- Pauli Exclusion Principle - Maximum of two electrons per suborbital (), each having an opposite spin.
- Hund’s Rule - Each suborbital () must have an electron of a certain spin before they can start refilling each for a maximum of two elections each. (Share the Cookie Rule).
- Aufbau Principle – electrons must fill lowest energy levels first before moving to next higher level
- Octect Rule - Completed valence electron energy level
B. Know at least: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2
C. Groups/Families are columns and have same # valence e- for A (Primary) elements
1. 10 elements have symbols that don’t match the first letters of name
2. Hg and Br only two natural liquids – all others solid or gas
D. Metals in A, and those in B (Secondary) are Transition Metals
1. Metals are to left of stair line (includes Al)
2. Lanthanides and Actinides are Transition Metals (Period 6 and 7)
E. Nonmetals – to right of stairstep line (Halogens and Noble Gases)
F. Metalloids – have side along stair line (but Al a metal)
Scientists and Summaries
DemocritusFirst to believe matter made of atoms (atomos)
AristotleMatter made of four elements (fire, air, earth and water)
Used by Church for nearly 1500 yrs.
R. BoyleExperiment to gain knowledge
Gases made of corpuscles (atoms) w/ spaces
J. DaltonAtomic Theory
-All elements made of atoms
-All elements of same kind have same atoms
-Reactions are changing, etc. of atoms
Conservation of Matter
W. CrookesUsed CRT to find electrons
JJ ThompsonDetermined charge on electron to be negative(magnets on CRT)
Plum Pudding Model of Atom
R. MillikanMass and charge amount of electron
Oil Drop Experiment
E. RutherfordGold Foil Experiment
Used alpha radiation (a He nucleus) to find protons
Beta and gamma radiation
J. ChadwickFound neutron (no charge)
N. BohrAtoms give off certain colors/wavelengths of light
Planetary Model of Atom
EinsteinBrownian Motion – atoms shown to exist by hitting pollen
Photoelectric Effect – light is wave or photon
Light (photon) is a packet of energy w/ no mass
E. SchrodingerElectrons at speed of light
Electron Cloud Model
W. HeisenbergUncertainty Principle – never certain of election place, use Probability to predict
Light used to see electrons makes them move
P. DiracAnti-matter
L. De BroglieElectron a wave and particle too
Four Forces
StrongHolds nucleus together with gluons
WeakHold each neutron and proton together (quarks) w/
bosons
EMFHolds nucleus to electrons - photons
GravityNo one knows – maybe a graviton
Standard Model of AtomIncludes subatomic particles (260+)
Developed from Quantum Theory
All matter made of Fermion particles
Baryonic Matter made of quarks (Hadrons -3; Mesons – 2)
Non-Baryonic includes leptons (electrons)
QuarksMake up Baryons
3 in proton and 3 in neutron–known as Hadrons
2 in Mesons
6 Types – up, down, top, bottom, charm, strange
I. Introduction - Chemical Bonds and Their Names
A. Bonds
1. Ionic, Covalent and Metal (alloys)
a. Main A Elements – use Roman Numeral to determine number of
valence e-‘s
(1) Mono, Di, Tri, Tetravalent set ups
(2) One pair of bonding e- (BP – bonding pair) per side
(3) Gaps show where bonds can occur, full side is Lone Pair (LP)
2. Remember – an ion is an element that has lost/gained an e- (‘s)
a. Must adjust diagram accordingly
(1) Cation (+ ion) – loses an e-; tend to be metals
(2) Anion (- ion) – gained an e-; tend to be non-metals
II. Bonds In General
A. Chemical Bond – occurs when e-‘s are simultaneously attracted to 2+ nuclei
1. Bonds can be diatomic or polyatomic
2. Based on EMF (plus charge to negative); or sharing electrons
3. Group VIII – full octet, no bonds
B. Pure / Non-Polar Covalent Bonds – e- pair is equally shared / attracted
between 2 nuclei
C. Polar Covalent – partial charges due to slight inequality in sharing
D. Ionic Bond – e- lost / gained due to electrostatic attraction
E. Metallic Bond – creates delocalized e- and sea of e- model
1. Alloys – two metals mixed
III. Ionic Bonds
A. Electrostatic force holding oppositely charged particles together
B. Metal cation w/ Non-metal anion
1. Binary and Ternary
a. Monatomic, Diatomic and Polyatomic ions
b. Oxyanions
2. Oxidation numbers
a. Criss-Cross Rule
C. Formation is always exothermic
1. Form crystal lattice structures in about 10 variations – salts
2. High melt / boil points
3. Hard and brittle
4. Conduct electricity when dissolved in water
a. Electrolyte
IV. Metallic Bonds
A. Metals share some properties w/ ionic compounds
B. Form lattice structures
1. Share valence e-‘s with surrounding atoms
2. Creates Electron Sea Model where e-‘s can shift around from one
nuclei to another
a. Delocalized e-
b. Makes metallic cation
c. Permits conductivity
V. Name and Formulas for Ionic Compounds
A. Formulas show simplest ratio of ions represented in ionic compound
1. Called formula unit
2. ex. Kbr = 1:1 ratio of potassium and bromine
MgCl2 = 1:2 ratio of magnesium to 2 chlorine ions
3. Overall charge of formula unit must = 0
a. ex. Mg2+ ion + 2 Cl- ions to form Mg Cl2
b. Overall charge of unit must = 0
4. Binary ionic compounds are metal cation and non-metal anion
a. Oxidation number important - # of transferable e-‘s
B. Naming Ionic Compounds
1. Binary Compounds
a. 1st word is name of the metal cation which remains same as
element
b. 2nd word is name of nonmetal anion
(1) If polyatomic – best to look up on lists
(2) If monatomic – replace ending of anion with “ide”
c. Determine if Roman Numeral needed to indicate oxidation
number of cation and place it in ( ) between words
(1) ex. FeCl2 and FeCl3 are both iron chloride but are different molecules
(2) Nickel Sulfide is prime example. Ni has +2, +3, and
+4 oxidation numbers and therefore must select correct
Roman Numeral for compound name. Since sulfide is -2,
use Nickel(II)Sulfide
(a) Transition Metals often have more than one
+ oxidation numbers
2. Ternary Compounds
a. Write name of metal cation first
b. Write name of polyatomic ion next
(1) Check for proper oxidation numbers
c. ex. Ca(CN)2 is calcium cyanide
Fe(NO3)2 is iron (II) nitrate; the last 2 only applies to
the NO3
3. Oxyanions – an ion with oxygen attached; most are polyatomic
a. Ion w/ more O atoms named using its root plus suffix of “ate”
b. Ion w/ fewer O atoms, use nonmetal root and add suffix “ite”
c. ex. Cl can form 4 oxyanions and are named according to # O
atoms
(1) Oxyanion w/ most O, use name w/ prefix of “per” and
suffix “ate”
(2) The one w/ 1 less, named w/ root of nonmetal and “ate”
suffix
(3) 2 fewer – root plus suffix “ite”
(4) 3 fewer O atoms – use root name of nonmetal with
prefix “hypo” and suffix “ite”
ex.: ClO4-: perchlorate ClO3-: chlorate ClO2-: chlorite ClO-: hypochlorite
4. Polyatomic Ions – 2+ atoms joined together that act as a single entity
a. Use lists and learn them!
b. Write formulas and names using binary rules
c. Watch oxidation numbers and use ( ) if needed
5. Metallic Ions
a. Many use original Latin names as root for metal cations
b. Lowest oxidations number has suffix “ous” on Latin name
c. Next highest, suffix of “ic” on Latin root
d. Know these Latin names!
6. Remember: Groups I, II, III A form only one ion
a. Nonmetals – take Group # and subtract 8 to find ion charge
b. Determine names from formula units and vice versa
(1) Criss-Cross Rule for oxidation numbers
VI. Covalent Bond Properties
A. Comparison of Ionics with Covalents
IonicsCovalents
Form crystalline solids – crystal latticeForm g, l, and/or s
(salts)
Give/Take 1+ electrons – from M to NMShare 1+ PAIRS of electrons
Exothermic ReactionsEndothermic reactions (mainly)
High melting/boiling points due to strongLow melt/boil points since atoms
Bonds (e- transfer)remain somewhat independent
Conduct thermal/electrical energy (e- canPoor conduction, insulators - no
be transferred or lost)“free” e-‘s to leave/move around
Hard, brittleSoft, malleable, flammable
Soluble in H2O (like dissolves like)Nonsoluble in nonpolar solutions
Nonsoluble in nonpolar liquidsSoluble in polar liquids like H2O
Metal cation to Nonmetal anion bondNonmetal to Nonmetal bond
VII. Naming Covalent Compounds
- Use prefixes for both nonmetals
- Mono (1) not always used for first element
- 2-di, 3-tr-, 4-tetra, 5-penta, 6-hexa, 7-hepta, 8-octa, 9-nona, 10-deca
- Remember diatomics
- H, N, O, F, Cl, Br, I (come in pairs if by themselves)
- S8 and P4
VIII. Chemical Reactions (Rxns)
A. Definition
B. Law of Conservation of Mass
C. Reaction mechanisms
1. Reactants
2. Products
3. Physical states – g, l, s, aq
4. indicates change (Δ) due to some process
D. Equation and Meanings
1. 2Al + 3 FeO Al2O3 + 2Fe
a. 2 atoms of Al + 3 molecules FeO in a reaction yield/make 1
molecule Al2O3 + 2 atoms Fe
b. 2 mols Al + 3 mols FeO yield 1 mol Al2O3 and 2 mols Fe
c. Using the Conservation of Mass:
(1) 2 x M of Al + 3 x M of FeO yields 1 x M Al2O3 + 2 x M Fe
(2) Molar mass of reactants = Molar mass of products –
regardless of their physical states and changes in these
2. Coefficients (stoichiometric coefficients) go in from of an atom or
molecule; subscripts are after an atom or molecule
IX. Balancing Chemical Equations
A. Identify all reactants and products
1. Place these in a skeleton equation (unbalanced form)
a. Make sure all molecules are balanced chargewise using the criss-cross
rule on charges if ionic, and prefixes if covalent
B. Make a list under the reactants and products of all elements involved and the number of atoms present
1. Be careful of applying subscripts appropriately
a. NH3 means 1 nitrogen and 3 hydrogen; whereas