Download Ionic Compounds and Their Names

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
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
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Study Guide – Fall Final
Physical Science – Dr. Hazlett
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 (lizard 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
2. Liquid to Solid
=
freezing
3. Liquid to Vapor
=
evaporation
4. Vapor to Liquid
=
condensation
5. Solid to Vapor
=
sublimation
6. Vapor to Solid
=
deposition
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Study Guide – Fall Final
Physical Science – Dr. Hazlett
The Atom and Periodic Table
I. The Atom
A. Nucleus
1. p+  gluon (strong force)  no
2. Each p+ and n0 made 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 ebefore any one gets a second eIII. 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 - s
b. B groups/Transition Metals - d
c. Lanthanides and Actinides - f
d. Groups III through VIII - p
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Study Guide – Fall Final
Physical Science – Dr. Hazlett
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 eB. 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
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Study Guide – Fall Final
Physical Science – Dr. Hazlett
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
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1. Both ideas fall apart due to exceptions/new elements discovered
D. Kekule and Meyer
1. Valency
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 eValence Energy Level (n)
1
2
3
4
5
6
7
Valence Shapes
s
s, p
s, p, d
s, p, d, f
s, p, d, f
s, p, d, f
s, p, d, f
Max. Val. e2
2, 6 = 8
2, 6, 10 = 18
2, 6, 10, 14 = 32
2, 6, 10, 14 = 32
2, 6, 10, 14 = 32
2, 6, 10, 14 = 32
1.
Shapes
a. S is sphere
b. P is principle or a figure 8
c. D is diffuse – four leaf clove
d. F is fundamental or 3d six leaf clover
2. Pauli Exclusion Principle each having an opposite spin.
3. Hund’s Rule (Share the Cookie Rule).
4. Aufbau Principle – electrons must fill lowest energy levels first before
moving to next higher level
5. 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 stair line (Halogens and Noble Gases)
F. Metalloids – have side along stair line (Al a metal)
Study Guide – Fall Final
Physical Science – Dr. Hazlett
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Scientists and Summaries
Democritus
First to believe matter made of atoms (atomos)
Aristotle
Matter made of four elements (fire, air, earth and water)
Used by Church for nearly 1500 yrs.
Experiment to gain knowledge
Gases made of corpuscles (atoms) w/ spaces
J. Dalton
Atomic Theory
-All elements made of atoms
-All elements of same kind have same atoms
-Reactions are changing, etc. of atoms
Conservation of Matter
W. Crookes
Used CRT to find electrons
JJ Thompson
Determined charge on electron to be negative (magnets
on CRT)
Plum Pudding Model of Atom
R. Millikan
Mass and charge amount of electron
Oil Drop Experiment
E. Rutherford
Gold Foil Experiment
Used alpha radiation (a He nucleus) to find protons
Beta and gamma radiation
J. Chadwick
Found neutron (no charge)
N. Bohr
Atoms give off certain colors/wavelengths of light
Planetary Model of Atom
Einstein
Brownian 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. Schrodinger
Electrons at speed of light
Electron Cloud Model
W. Heisenberg
Uncertainty Principle – never certain of election
place
Light used to see electrons makes them move
P. Dirac
Anti-matter
L. De Broglie
Electron a wave and particle too
Study Guide – Fall Final
Physical Science – Dr. Hazlett
Four Forces
Strong
Weak
EMF
Gravity
8
Holds nucleus together with gluons
Hold each neutron and proton together (quarks) w/
bosons
Holds nucleus to electrons - photons
No one knows – maybe a graviton
Standard Model of Atom
Includes 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)
Quarks
Make up Baryons
3 in proton and 3 in neutron - 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
Study Guide – Fall Final
Physical Science – Dr. Hazlett
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 eb. 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
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Study Guide – Fall Final
Physical Science – Dr. Hazlett
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2. Ternary Compounds
a. Write name of metal cation first
b. Write name of polyatomic ion next
(1) Check for proper oxidation numbers
ex.
c.
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
Ionics
Form crystalline solids – crystal lattice
(salts)
Covalents
Form g, l, and/or s
Give/Take 1+ electrons – from M to NM
Share 1+ PAIRS of electrons
Exothermic Reactions
Endothermic reactions (mainly)
High melting/boiling points due to strong
Bonds (e- transfer)
Low melt/boil points since atoms
remain somewhat independent
Study Guide – Fall Final
Physical Science – Dr. Hazlett
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Conduct thermal/electrical energy (e- can
be transferred or lost)
Poor conduction, insulators - no
“free” e-‘s to leave/move around
Hard, brittle
Soft, malleable, flammable
Soluble in H2O (like dissolves like)
Nonsoluble in nonpolar solutions
Nonsoluble in nonpolar liquids
Soluble in polar liquids like H2O
Metal cation to Nonmetal anion bond
Nonmetal to Nonmetal bond
VII. Naming Covalent Compounds
A. Use prefixes for both nonmetals
a. Mono (1) not always used for first element
b. 2-di, 3-tr-, 4-tetra, 5-penta, 6-hexa, 7-hepta, 8-octa, 9-nona, 10-deca
B. Remember diatomics
a. H, N, O, F, Cl, Br, I (come in pairs if by themselves)
b. 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
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Physical Science – Dr. Hazlett
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a. NH3 means 1 nitrogen and 3 hydrogen; whereas
(NH3)2 means the outer 2 is multiplied toward the inside subscripts giving
2 nitrogens and 6 hydrogens
C. Start to balance both sides of equation by adjusting coefficients (front number)
1. Do not place a coefficient inside a molecule – only in front – since
doing so changes the entire structure of the molecule
a. Coefficients multiply out to all members’ subscripts of a
molecule
(1) 2(NH3) means 2 nitrogens and 6 hydrogens; and
2(NH3)2 means the 2 nitrogens from the subscript are now
multiplied by 2 to make 4, and the 3 hydrogen are multiplied by 2 by the
outer subscript and then 2 again by the coefficient making 12 total
2. Do not change, insert or modify existing subscripts for any reason!
D. Change the number of atoms via the coefficients used until both sides
are equal and balanced
E. Hints:
1. List atoms under equation in same order to avoid confusion
2. Keep atoms of same element but in different molecules “separate” so can
adjust accordingly if needed
a. ex. H + H2O ; the H are listed as 1 + 2 and can multiply
one or the other as needed
3. Try doing H and O last if possible
4. Remember the diatomics when writing an equations, and S8, P4
5. Reduce a final equation if all coefficients have a common divisor
X. Why Rxns Occur?
A. Collision Theory
1. Chemical Kinetics
a. Molecules in a g, l are in constant random motion – KE
(1) Rotational, Translational, Vibrational motion
(2) Collisions occur due to KE
B. Concentration of Reactants
1. Higher the concentration – more collisions and more rxn
C. Temperature and Pressure
1. Increase in T and/or P increases KE and collisions, thus rxns
D. Surface Area
1. More surface area – more collisions
a. Called adsorption
E. Catalysts
1. Used to speed up or slow down reactions especially with acids/bases
a. These called buffers
2. Adsorption – reacting molecules bond to surface of catalyst
a. Catalysts provide a new pathway or mechanism for the
molecule to react, i.e. a change in the process
F. Agitation
1. Includes mixing, shaking, stirring the reactants to create products
2. If solid is a product – called precipitate
I. Activation Energy (Ea)
1. Svante Arrhenius (1888) – temperature affects the collision of
Study Guide – Fall Final
Physical Science – Dr. Hazlett
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molecules and gives the activation energy needed to break (Do – dissociation
energy) or form bonds, making rxns
XI. Types of Reactions
1. Synthesis reaction: occurs when two or more simple substances (elements or
compounds) are combined to form one new and more complex substance. The general
form of a synthesis reaction is:
 element or compound + element or compound compound
 Fe + S FeS
a. The driving force for these reactions is always the transfer of electrons from the
less electronegative element to the more electronegative element.
b. Nonmetal + nonmetal  Covalent compound
c. Metal + Nonmetal  Ionic Compound
2.
Single displacement reaction: occurs when one element displaces another in a
compound. The general form of a single displacement reaction is:
 element + compound element + compound
 Zn + 2HCl H2 + ZnCl2
3.
Double displacement reaction: occurs when the cation (+) and the anion (-) of
the two reactants are interchanged. The general form of a double displacement
reaction is:
 compound (AB) + compound (CD) compound (AD) + compound
(CB)
 FeS +2HCl FeCl2 + H2S
4.
Decomposition reaction: occurs when energy in the form of heat, light,
electricity, or mechanical shock is supplied. A compound may decompose to
form simpler compounds and/or elements. The general form of a decomposition
reaction is:
 compound two or more substances
5. Combustion: A combustion reaction is when oxygen combines with another
compound to form water and carbon dioxide. These reactions are exothermic, meaning
they produce heat. C10H8 + 12 O2  10 CO2 + 4 H2O
6. Acid-base: This is a special kind of double displacement reaction that takes place
when an acid and base react with each other. The H+ ion in the acid reacts with the OHion in the base, causing the formation of water. Generally, the product of this reaction is
some ionic salt and water:



HA + BOH  H2O + BA
2NaOH(aq) + H2SO4(aq)  Na2SO4(aq) + 2H2O(l)
One example of an acid-base reaction is the reaction of hydrobromic acid (HBr)
with sodium hydroxide: HBr + NaOH ---> NaBr + H2O
Study Guide – Fall Final
Physical Science – Dr. Hazlett
XII.
Solutions
A. Homogeneous v. Heterogeneous Mixtures
B. Soln = Solvent + Solute
C. Solubility – likes dissolve likes (polarity)
1. Miscible
2. Solvation Process
3. Saturation
4. Colligative Properties – Lowering Vapor pressure;
Boil/Freeze Temp Changes; Osmotic Pressure
5. Dilution and Concentration
XIII. Introduction to Acids and Bases
A. Acids (A) and Bases (B) are in solutions (solns)
1. Solns are homogeneous mixtures of 2+ substances
a. Solute and solvent
2. Electrolytes and non-electrolytes
a. Aqueous (aq) solns of ionic compounds conduct electricity – i.e.
flow of electrons, are electrolytes
b. Pure (distilled) water not a good conductor – no ionization and
therefore, no free electrons to conduct electrical energy
c. A’s ionize in water – thus are electrolytes
(1) Ionization varies up to 100%
(2) Weaker acids ionize less than 100%
3. Ionization
a. Aqueous soln (water)  end up with H+ and OH- ions
b. The amounts of these ions in soln determine whether A or B
c. Water is universal solvent
(1) Undergoes self-ionization
(a) H2O  H+ + OH(b) H2O + H2O  H3O+ + OHXIV. Acids and Bases in General
A. Characteristics
1. Acids - Physical:
a. Taste sour
b. React with some metals (Al, Mg, Zn) to produce H2 (g)
c. React with carbonates to produce CO2 (g)
d. React with B to form salts and water
e. Change indicators to differing colors – most common is
litmus blue to red
f. Yields H+ ion or forms H3O+
2. Bases – Physical:
a. Taste bitter
b. Fell soapy/slippery
c. Changes red litmus paper to blue
d. Reacts with A’s to form salts and water (neutralization)
e. Reacts with some metals to produce H2 (g)
f. Yields the OH- ion
g. Alkalines
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Study Guide – Fall Final
Physical Science – Dr. Hazlett
15
B. Strong Acids/Bases
1. Strong Acid – completely ionize/dissociate in aqueous soln
a. Only 7: HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4
(1) Sometimes H3O+, HBrO3, HBrO4, HIO3, and HIO4 included
b. Weak Acids – do not ionize/dissociate completely
c. Strength depends on polarity of bond between H and the element(s) to
which it is bonded
(1) Connected to the level of Do
(2) Strength increases when polarity increases with decrease in
bond energy
2. Superacids – pH way beyond 14 – up to 2 x 1019 stronger than H2SO4
3. Strong Base – completely ionize/dissociate in soln
a. Only 7: LiOH, NaOH, KOH, RbOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
b. Weak bases do not ionize completely
4. Salts from AB Rxns
a. Salt is an ionic cmpd made from cation of a B and an anion of an A
(1) Metal with a nonmetal  salt
XV. The Mole
A. KNOW THIS PROCESS!
÷ Molar Mass (M)
-------------------------
x NA
----------------------
Mass
Moles
Number of
(Grams)
n
Particles**
--------------------------------------------x Molar Mass (M)
÷ NA
* NA = 6.022314 x 1023
** Particles = ions, atoms, molecules
B. NA is Avogardro’s Number and equals 6.022 x 1023
1. This became equated with the “mole” (abbreviated mol)
a. 1 mole of anything = 6.022 x 1023 of that item
C. Molar Mass (M) is the number of grams per mole of a substance (g/mol)
1. For elements on Periodic Chart – look for atomic mass (amu)
under symbol and change this to xx.xxx g/mol
2. For compounds – add up all masses of all elements in the compound
a. Careful to watch for subscripts and distributive property
b. For compounds – this is called Molecular Molar Mass
3. Example (some of these are rounded off):
Study Guide – Fall Final
Physical Science – Dr. Hazlett
Element/Compound
C
He
K
H2O
C6H12O6
16
Atomic/Molecular Mass
Molar Mass
12.011 amu (u)
12.011 g
4.00 u
4.00 g
39.098 u
39.098 g
(2 x 1.007) + 15.999 = 18 u
18 g
(6 x 12) + (12 x 1) + (6 x 16) = 180 u 180 g
D. Moles (mols or n)
1. An encompassing term for 6.022 x 1023 units of whatever is being
counted, regardless of their mass
a. Therefore, 1 mol of whatever units/particles = 6.022 x 1023
b. Derived from C-12 atom and its atomic mass
2. # mols = # particles ÷ NA
3. # mols = mass ÷ molar mass
4. # mols = concentration (Molar) ÷ volume (Liter)
5. Mass = # mols x molar mass
6. # particles = # mols x NA
7. Atomic mass = mass of element ÷ # mols
XVI. Periodic Trends (Honors Only)
I. Effective nuclear charge – the attractive positive charge of nuclear protons acting on valence
electrons.
1. The effective nuclear charge is always less than the total number of protons present in a nucleus
due to shielding effect.
2. Effective nuclear charge is behind all other periodic table tendencies.
II. Shielding effect – the lessening of attractive electrostatic charge difference between nuclear protons
and valence electrons by partially or fully filled inner shells.
A. Shielding effect increases with the number of inner shells of electrons.
1. Electrons sharing the same shell do not shield one another from the attractive pull of the nucleus.
B. The periodic table tendency for effective nuclear charge:
1. Increase across a period (due to increasing nuclear charge with no accompanying increase in
shielding effect).
2. Decrease down a group (although nuclear charge increases down a group, shielding effect more
than counters its effect).
C. Effective nuclear charge can be used to explain the octet rule.
1. The octet rule is the very strong tendency of an element to adopt the electron configuration of
the Noble gas (numerically) closest to it in the periodic table.
III. Atomic Radii - atomic radius an estimate of the size of an atom.
A. The periodic table tendency for atomic radius:
1. Decrease across a period (due to increasing nuclear charge with no accompanying increase in
shielding effect).
2. Increase down a group (although nuclear charge increases down a group, shielding effect more
than counters its effect).
IV. Ionic Radii - ionic radius is an estimate of the size of an ion.
A. The periodic table tendency for ionic radius:
Study Guide – Fall Final
Physical Science – Dr. Hazlett
17
1. Cations are smaller than their atoms. There are two effects at work:
The loss of valence electrons results in the loss of the outermost shell.
The remaining electrons actually experience an increase in nuclear charge
with no accompanying increase in shielding effect.
2. Anions are larger than their atoms (additional electrons with no accompanying change in either
nuclear charge or shielding effect).
3. Increase down a group for ions carrying the same charge.
4. Decrease across an isoelectronic series.
An isoelectronic series is a series of ions all of whom have the same electron configuration.
For example; O-2, F-1, Na+1, Mg+2, and Al+3 all have the same electron configuration, namely, [Ne].
V. Ionization Energy - The ionization energy of an atom or ion is the minimum amount of energy
necessary to remove an electron from an isolated gaseous atom or a ground state ion.
A. Every electron in an atom or ion has a characteristic ionization energy, and in general,
1. Valence electrons have lower ionization than core electrons.
2. First ionization energies are lower than successive ionization energies.
B. The periodic table tendency for first ionization energy:
1. Generally increase across a period (due to increasing nuclear charge with no accompanying
increase in shielding effect).
VI. Electron Affinity - The energy change that occurs when a gaseous atom or ground state ion gains
an electron is termed electron affinity.
A. The periodic table tendency for electron affinity:
1. Generally becomes increasingly negative across a period.
-It is important to keep in mind that the more negative the value for electron affinity becomes, the
greater the attraction for an additional electron.
-It is therefore accurate to state that electron affinity generally increases across a period.
2. Generally becomes decreasingly negative down a group, although this is a very small effect in most
groups.
VII. Metallic Character - The stair-step line divides the periodic table into metals, nonmetals, and
metalloids.
A. 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.
B. Properties of metals include:
1. Shiny luster
2. Malleability (can be rolled into thin sheets)
3. High melting point
4. Form ionic compounds with nonmetals
D. The periodic table tendency for metallic character:
1. Generally decreases across a period.
2. Generally increases down a group.
5. Conductivity of heat and electricity
6. Ductility (can be pulled into thin wires)
7. Low first ionization energy
8. Form basic oxides