Chemistry
Answers to Explain (Analysis), Part 1:
1. Examine the data collected for melting point. What conclusions can you draw about the
melting point of these chemicals?
- The chemicals that took longer to melt have a higher
melting point than those that melted more quickly.
2. Which substances have higher melting points?
Which have lower melting points? What
does this indicate about the bonds in the substances?
-Substances with ionic bonds had higher melting points
and those with covalent bonds had lower melting points.
A lower melting point indicates weaker bonds that will be
more easily broken. A higher melting point indicates
stronger bonds.
3. Summarize the solubility of the substances in the Explore Activity.
-All substances are soluble in water, except benzoic acid.
4. How is solubility associated with the type of bond present?
-All our ionic substances are soluble in water. However,
not all ionic compounds are soluble in water. One of
our covalent molecules are soluble (dextrose) and one is
not (benzoic acid).
5. What does the solubility of the different substances indicate about the type of bond
present?
-Substances with weaker intermolecular forces are more
soluble in water. Those with stronger intermolecular
forces are less soluble in water.
6. How is conductivity related to the type of bond present?
-Ionic substances, when dissolved in water, conduct
electricity. Covalent substances do not conduct
electricity when dissolved in water.
7. Why do substances with certain types of bonds conduct electricity well, while some
substances are not good conductors?
-Ionic compounds conduct electricity well because
they possess ions, which allows electrons to flow from
atom to atom. This occurs only when they are melted
or dissolved in water.
8. Is there a significant difference in appearance between the substances with covalent
bonds and those with ionic bonds? What properties did you notice you could not see
with the naked eye?
-Both ionic and covalent compounds appear to be
white solids. Under the hand lens, however, the
covalent substances are smaller in particle size than
the particles in ionic compounds. Ionic particles also
have a more geometric, crystalline shape while
covalent particles vary in shape.
9. Imagine looking at the substances under a microscope. What do you think the substances
might look like on a microscopic level?
-(Answers will vary.)
Lewis Structures are used to show bonding in molecules
and ionic compounds.
-dots represent valence electrons
-in ionic bonding, the charge of each ion must be shown
-in covalent bonding, bonded electrons is shown by
lines
Arrows represent transfer of electrons from the metal to
the nonmetal.
-the charge of each atom must be shown
Example: CaS
When we show bonding, shared electron pairs can be
shown by either a pair of dots or a single line.
-Lewis Structures are used to show how bonding
electrons are arranged in molecules
-example: NH3
-sigma bond (s): single covalent bond formed when
an electron pair is shared by the direct overlap of
orbitals
♦can occur between s & s, s & p , or p & p orbitals
Explain(Analysis), Part 2:
Draw Lewis structures for the following ionic
compounds.
1. NaCl
2. MgO
3. LiF
Draw Lewis structures for the following covalent
compounds.
1. H2O
2. CO2
3. NH4
A multiple bond forms when two atoms share more than 2
electrons.
-double bond: 4 electrons shared ( 2 pairs)
♦ O2
-triple bond: 6 electrons shared (3 pairs)
♦ N2
Some molecules have both single and multiple bonds.
♦HCN
pi bond (p): forms when parallel orbitals overlap to share
electrons
-only occurs with multiple bonds because the first
overlap is always a sigma bond
Show the formation of the ionic compound for the
following pairs of elements
1. strontium and fluorine
2. aluminum and oxygen
3. cesium and phosphorus
4. lead and chlorine
5. potassium and iodine
6. magnesium and chloride
7. aluminum and bromide
8. cesium and nitride
9. barium and sulfide
1. PH3
2. H2S
3. HCl
4. SCl2
5. SiH4
6. CO2
7. CH2O
8. C2H2
chemical bond: force that holds two atoms together
-creates stability in the atom
Two types of bonds:
1. Attraction between a positive nucleus and negative
electrons (covalent bonding)
2. Attraction between a positive ion and a negative ion
(ionic bonding)
Remember: It is the valence electrons that are involved in
this bonding.
ionic bond: electrostatic force that holds oppositely
charged particles together
-called ionic compounds
-forms between metals and nonmetals
◊metals lose electrons, forms a cation
~cation: positive ion from loss of electrons
◊nonmetals gain electrons, forms an anion
~anion: negative ion formed from gain of electrons
-most are binary, which means they contain 2 different
elements, such as MgO, Al2O3
It is the chemical bonds between atoms that determines
many of the physical properties of the compound.
-alternating positive and negative ions form an ionic
crystal
-the ratio of positive to negative ions is determined
by the number of electrons transferred
-strong attraction
results in a crystal
lattice, a 3-D
arrangement of
atoms.
Other characteristics include:
-high melting and boiling points
-very hard and rigid
-brittle
-electrolyte when dissolved in water (aqueous solution)
During chemical reactions, energy is either absorbed
(endothermic) or released (exothermic)
-the formation of ionic bonds is always exothermic
lattice energy: energy required to separate one mole
of ions of an ionic compound
-the more negative the lattice energy, the stronger the
bond
Lattice Energies of Some Ionic Compounds
Compound
Name
Lattice Energy
(kJ/mol)
Compound
Name
Lattice Energy
(kJ/mol)
KI
-632
KF
-808
KBr
-671
AgCl
-910
RbF
-774
NaF
-910
NaI
-682
LiF
-1030
NaBr
-732
SrCl2
-2142
NaCl
-769
MgO
-3795
Compound
Name
KI
KBr
RbF
NaI
NaBr
NaCl
Lattice Energyies of Some Ionic Compounds
Lattice Energy
Compound
Lattice Energy
(kJ/mol)
Name
(kJ/mol)
-632
KF
-808
-671
AgCl
-910
-774
NaF
-910
-682
LiF
-1030
-732
SrCl2
-2142
-769
MgO
-3795
Depends on:
1. smaller ions -more negative value because the attraction is
stronger between the nucleus and valence electrons
2. larger the positive/negative charge, the more
negative the lattice energy because the attraction is
stronger when more electrons are lost/gained
1. Draw the Lewis dot notation showing the bonding
between beryllium and chlorine.
2. What determines the properties of an element?
3. What is a crystal lattice?
4. List 5 characteristics of ionic compounds.
5. What is the difference between endothermic and
exothermic? Which occurs in ionic reactions?
6. What is lattice energy?
7. What does lattice energy depend on?
8. Which substance has a stronger bond: NaCl or NaBr?
Why?
Remember that atoms bond to increase stability, which
occurs when an atom gets a full outer shell of electrons.
-in ionic bonding, one atom loses electrons (metal)
and another gains electrons (nonmetal) to form
oppositely charged ions with a full outer shell
However, sometimes there is not a transfer of electrons,
but a sharing of electrons.
-covalent bond: attractive force between atoms due
to the sharing of valence electrons
Covalent bonds can form between:
-2 or more nonmetal atoms
-metalloids (especially the ones to the right of the
metalloid line) and nonmetals
molecule: when two or more atoms bond covalently
Covalent bonds can have either single bonds or multiple
bonds.
-single bonds: 2 shared electrons (1 pair)
-multiple bonds: 4 or 6 electrons shared (2 pair=
double or 3 pair = triple)
1. low melting and boiling points.
2. many vaporize readily at room temperature
3. relatively soft solids (but not all, some are gases/liq.)
4. can form weak crystal lattices
5. do not conduct electricity when dissolved in water
These properties are due as a result of differences in
attractive forces
-attraction between atoms within a molecules is strong
-attraction between different molecules is weak
~called intermolecular forces or van der Walls forces
Types of Intermolecular Forces (van der Walls forces)
1. dispersion force (induced dipole)
2. dipole-dipole force
3. hydrogen bonding
dispersion force (induced dipole)
-occurs between nonpolar molecules
-very weak
dipole-dipole force
-occurs between polar molecules
-the more polar the molecule, the stronger the force
hydrogen bonding
-strong intermolecular force between the hydrogen
end of one dipole and a fluorine, oxygen or nitrogen
atom on another molecule’s dipole
All bonds can be broken, though some more easily than
others.
-due to the strength of the bond
What affects bond strength?
bond length: distance that separates the bonded nuclei
-determined by the size of the atoms and how many
electron pairs are shared
♦larger the atom, the longer the bond length, the
weaker the bond
♦more shared electrons gives a shorter, stronger bond
When a bond forms or breaks, an energy change occurs.
-bond formation: energy released (exergonic)
-bond breaking: energy absorbed (endergonic)
bond dissociation energy: amount of energy required to
break a specific covalent bond
-always a positive number
-indicates the strength of a covalent bond
larger the bond dissociation energy, stronger the bond
(see p 246 for examples)
1.
2.
3.
4.
5.
6.
7.
8.
9.
Describe a covalent bond.
What types of atoms do covalent bonds form between?
Describe single, double and triple bonds.
What do we mean by sigma and pi bonds?
What do we call covalent compounds?
What affects bond strength?
Describe the two things that determine bond length.
What does bond dissociation energy indicate?
What occurs when a bond forms or breaks?
Remember that atoms have different attractions for
electrons (electronegativity).
-electronegativity increases left to right and decreases
down a period
The character and type of bond can be predicted using the
difference in electronegativities between bonded atoms.
-pure covalent bond: electronegativity difference = 0
(usually occurs between identical atoms, H2)
Most atoms do not have equal sharing of electrons,
producing a purely covalent bond.
-polar covalent bond: unequal sharing of electrons
♦the larger the electronegativity difference, the more
ionic the bond character
-ionic bonds form when the electronegativity
difference is > 1.7 and nonpolar covalent bonds form
when the difference is < 0.5
-the cutoff between polar covalent, nonpolar, and
ionic is sometimes inconsistent with experimental
data
Remember: bonding is not clearly ionic or covalent, with ionic
character increasing as the difference in electronegativity
increases.
Decide if the following pairs of atoms are polar covalent,
nonpolar covalent or ionic.
1.
N-H
3.04-2.20 = 0.84
polar covalent
2. C-Cl
2.55-3.16 = 0.61
polar covalent
3. S-Se
2.58-2.55 = 0.03
nonpolar covalent
When a polar bond forms the shared electrons are pulled
more strongly toward one atom.
-this creates partial charges at opposite ends of the
molecule, which is called a dipole
♦ d- indicates a partial negative
d+ indicates a partial positive
Polar molecule or not?
A molecule can have individual polar bonds, but make a
nonpolar molecule. How?
We look at the shape of the molecule.
Let’s look at H2O and CCl4.
O—H
C—Cl
dd+
d+
d1.24
0.61
both O-H and C-Cl have polar covalent bonds
One molecule is polar and the other is nonpolar? How do
we know?
We look at the shape of the molecule and the terminal
atoms.
-symmetric molecules like CCl4 are nonpolar because the
polar bonds cancel each other out.
CCl4
-asymmetric molecules like H2O are polar because the
polar bonds do not cancel each other out.
H2O
If water is polar, why will oil not dissolve in it?
Oil must be nonpolar because
A substance is only soluble (dissolvable) when combined
with a like molecule.
“Like Dissolves Like”
hydrophobic- “fear of water”
hydrophilic- “likes water”
1. What is electronegativity and what does it predict?
2. What is the difference between a nonpolar covalent
bond and a polar covalent bond?
3. What is a dipole and what indicates them?
4. Describe the electronegativity trend both across a
period and down a group.
5. Are the following bonds polar or nonpolar covalent?
a. H-Br
b. C-O
c. S-C
6. Describe the relationship between polarity and
solubility.
7. What do we mean by symmetric and asymmetric?
Final Bonding Questions:
1. Draw a table comparing the properties of ionic and
covalent bonds.
-leave room to add more properties (we will discuss
the table and add more to it)
2. What is a general definition of a bond?
3. What are the two types of bonds? Describe each.
4. What is the octet rule?
5. What do we mean by polar or nonpolar?
6. What is electronegativity? How do we use this in
bonding?
7. What are intermolecular forces?
1.
Ionic
high melting/boiling point
electrolyte in water
crystal lattice structure
hard, brittle solids
most dissolve in water
Covalent
low melting/boiling point
nonelectrolyte in water
some form weak crystal lattices
gases, liquids, relatively soft solids
some dissolve in water
many vaporize at room temp
2. A bond is a force holding two atoms together to
create stability in an atom
3. An ionic bond is an attraction of oppositely charged
ions due to a transfer of electrons from a metal atom
to a nonmetal atom. A covalent bond is the sharing
of electrons between nonmetals or nonmetals and
some metalloids.
The octet rule states that atoms are stable if they
have a full valence shell of electrons. For most
atoms, the number is 8, but the period 1 elements
are stable with 2.
5. A covalent bond is polar if there is an unequal
sharing of electrons due to the electronegativity
difference between the atoms. It is nonpolar if there
is an equal sharing of electrons.
6. Electronegativity is the attraction an atom has for
electrons. The more electronegative the atom, the
stronger the attraction. We use electronegativity to
determine the polarity of molecules.
7. Intermolecular forces are the force that holds atoms
together. They can be weak, allowing atoms to be
pulled apart easily, or strong.
4.
TEST #1
structural formula: uses letter symbols and bonds to
show relative positions of atoms
-one of the most useful
-can be predicted for many molecules by drawing
Lewis structures
-H is always an end (terminal) atom, never a central
atom
-less electronegative atom is the central atom
(nm or metalloid closest to the left of the PT-usually)
CH2O
1. Predict the location of the atoms
C is least electronegative & farthest to left on PT,
therefore it is the central atom
2. Find the total number of electrons available for
bonding.
1 C-4, 2 H-2, 1 O-6 for a total of 12 valence e3. Determine the number of bonding pairs
12 valence e- / 2 = 6 electron pairs
4. Place one bonding pair (single bond) between the
central atom and each terminal atom.
H
C
O
H
5. Subtract the number of pairs you used in step 4 from
the number of bonding pairs determined in step 3.
6 – 3 used = 3 e- pairs left
5. Subtract the number of pairs you used in step 4 from
the number of bonding pairs determined in step 3.
-take the remaining electron pairs and place electron
pairs around the terminal atoms to satisfy the octet
rule
H
C O
H
6. If the central atom is not surrounded by 4 electron
pairs, it does not have an octet
-convert one or two of the lone pairs on a terminal
atom to a double or triple bond between that terminal
atom and the central atom
H C O
H
Practice:
1. CH3Cl
2. NBr5
Writing structural formulas for polyatomic ions is the
same with one exception:
-the total number of electrons may differ due to the
negative and positive charge.
♦negative charge, more electrons are present
SO4-2
add two electrons
♦positive charge, less electrons are present
NH4+1 subtract one electron
Let’s look at CO3-2.
-when one or more valid Lewis structure can be written for
a molecule, resonance occurs
-let’s look at another resonance molecule/ion: NO3-1
-each molecule/ion that undergoes resonance behaves as
if it only has one Lewis structure
Some molecules do not obey the octet rule.
Three reasons exist:
1. when a small group of molecules have an odd number of
valence electrons:
-NO2 for a total of 17 valance electrons-one unpaired
electron on N
2. Some form with fewer than eight, though this is
relatively rare:
-B in BH3 is stable with six because it only has 3 valence
electrons.
3. When the central atom has more than 8 electrons,
which is referred to as an expanded octet.
-can occur in elements that are found in period three
or higher elements (because of the d orbitals).
-P in PCl5
(1 s orbital, 3 p orbitals, and 1 d orbital)
1. SO3
2. N2O
3. SF6
4. ClF3
5. SiF4
6. PO4-3
7. BF3
8. SO3-2
1. What is a structural formula?
2. Describe resonance.
3. List three reasons for exceptions to the octet rule.
4. Name the following:
a. BH3
b. SO2
c. PO4-3
5. Write formulas for the following:
a. sulfur trioxide
c. chlorous acid
b. hydrosulfuric acid
6. Draw structural formulas
a. SO2
b. H2O
c. BrCl5
Many of the physical and chemical properties of molecules
is determined by the shape of the molecule.
-the shape of molecules determines if two or more
molecules can get close enough for a reaction to occur.
VSEPR (Valence Shell Electron Pair Repulsion)
model: atoms in a molecule are arranged so that the
pairs of electrons (bonded and lone) minimize
repulsion.
The repulsion between electron pairs result in fixed angles
between atoms
-bond angle: angle formed by any two terminal atoms
and the central atom
♦lone pairs take up slightly more space than bonded
pairs
♦multiple bonds have no affect on the geometry
because they exist in the same region as single
bonds
-example: H2O
See page 260 for the Molecular Geometries (Shapes)
1. What determines many of the physical and chemical properties of
molecules?
2. Describe the VSEPR model.
3. What does the repulsion between electron pairs result in?
4. Why do multiple bonds have no affect on geometry
of a molecule?
5. Why do molecules with lone pairs have shorter bond angles?
6. How do we know if a molecule is polar or nonpolar?
A universal set of rules must be used so chemists around
the world can communicate.
formula unit: simplest ratio of ions represented in an
ionic compound
-remember that ionic compounds form a crystal lattice,
consisting of many cations and anions.
-the overall charge for the compound is 0
Most ionic compounds are binary, consisting of two
monatomic ions.
-monatomic ion: one atom ion, either positively or
negatively charged
Remember that we determine the charge of each
ion by its oxidation number.
Formula Rules for Ionic Compounds
1. write the cation first, followed by the anion
2. state the charges of both ions
3. cross the number for the charge of one ion to become
the subscript for the other ion.
-subscripts are used to state the number of each atom
in the compound
Example: Determine the formula for the ionic compound
formed when potassium reacts with oxygen.
1. Cation = potassium = K
Anion = oxygen = O
2. K+1 O-2
3. K+1 O-2
K2O1
K2O
You try: Determine the formula for the ionic compound
formed when aluminum reacts with chlorine.
We write formulas for ionic compounds containing
polyatomic ions the same way as in binary compounds.
-the cation comes first, followed by the anion
-state the charges
-cross over the number for the charges
However:
-if you have more than one polyatomic ion, place
parenthesis around the polyatomic ion, with the
subscript outside the parenthesis.
Example: Determine the formula for the ionic compound
formed when beryllium reacts with cyanide.
1. Cation = beryllium = Be
Anion = cyanide = CN2. Be+2 CN-1
3. Be+2 CN-1
Be1(CN)2
Be(CN)2
You try: Determine the formula for the ionic compound
formed when ammonium reacts with iodine.
Write the correct formula for the following pairs of atoms:
1. ammonium and oxygen
2. lithium and nitrate
3. aluminum and hydroxide
4. ammonium and phosphate
5. strontium and acetate
1. Why do we need a universal set of rules for naming and
writing formulas?
2. Define monatomic and binary.
3. What is meant by a formula unit?
4. What is the purpose of subscripts.
5. Describe what a polyatomic ion is?
6. When do we use parenthesis for writing ionic compounds
with polyatomic ions?
7. Determine the formula for the ionic compound formed
when lead reacts with sulfur.
8. Determine the formula for the ionic compound formed
when lithium reacts with nitrogen.
Write the correct formula for the following pairs of atoms:
1. aluminum and carbon
2. ammonium and carbonate
3. calcium and oxygen
4. aluminum and chromate
5. sodium and phosphate
6. potassium and hydrogen sulfate
7. magnesium and phosphorus
The names of ionic compounds include the ions of which
they are composed.
1. The element whose symbol appears first in the
formula also appears first in the name.
-this is always the positively charged ion, or metal
2. The name of the second ion follows, with its ending
changed to –ide for single atom ions.
Ex: What is the name of MgCl2?
magnesium chloride
Write the formula and the name.
1. Na2S
2. Ga2S3
3. CaSe
4. LiF
You follow the same rules when naming polyatomic ions as
when you have binary ionic compounds, however:
-you do not change the ending of the polyatomic ions,
even when they are the second atom.
Example:
Al2(SO4)3
aluminum (III) sulfate
Rule: You must state the charge of all metals not
included in groups 1 and 2 because many have
multiple charges.
Name the following compounds:
1. NaC2H3O2
2. CaCO3
3. KOH
4. Mg(NO3)2
5. Li2CrO4
6. Mg(NO2)2
7. AlPO4
*According to the previous rules, FeO and Fe2O3 would
both be named iron oxide,even though they are not the
same compound*
Since many transition metals can have more than one
charge, the name must show this. This is done using
roman numerals.
-FeO
is named iron (II) oxide because Fe has a +2
charge
-Fe2O3
is named iron (III) oxide because Fe has a
+3 charge
*The roman numeral states the charge of the metal*
Q: How do I know the iron in FeO has a +2 charge?
A: The oxide ion has a –2 charge, so the Fe must have
a +2 charge so the compound is overall neutral.
Q: How do I know the iron in Fe2O3 has a +3 charge?
A: There are three oxide ions with a –2 charge:
(3 ions)(-2 charge/ion) = a total of –6 charge
Since the overall charge must be neutral, the iron
must have a total charge of +6. Therefore:
(2 ions)(x charge/ion) = +6
x = +3
Write the formula given & the name of each compound.
1. FeCl3
2. Zn3P2
3. CuS
4. AuF
5. CuC2H3O2
6. AgHCO3
7. ZnSO4
8. Pb(CO3)2
1. What is the ending of the second atom changed to when
naming ionic compounds?
2. Write the name for (NH4)3P
3. Write the name for AlS.
4. Determine the formula for the ionic compound formed
when magnesium reacts with phosphate.
5. Determine the formula for the ionic compound formed
when magnesium reacts with phosphate.
Molecules are represented by both names and formulas.
Rules for Naming Binary Molecular Compounds
1. The first element in the formula is named first, using
the entire element name.
2. The second element in the formula is named using
the root of the element and adding the suffix –ide.
3. Prefixes are used to indicate the number of atoms of
each type that are present in the compound.
-exception: 1st element never uses the prefix mono-drop the final letter of the prefix if element name
begins with a vowel.
Prefixes you need to know:
# atoms
prefix
1
mono2
di3
tri4
tetra5
penta6
hexa7
hepta8
octa9
nona10
deca-
Name the compound P2O5, which is used as a drying and
dehydrating agent.
1st atom: P = phosphorus
2nd atom: O = oxygen = oxide
There are 2 phosphorus = diphosphorus
There are 5 oxygens = pentoxide (drop the –a of penta-)
Put it together: diphosphorus pentoxide
Name the following molecules:
1. CCl4
2. As2O3
3. CO
4. SO2
5. NF3
(We will talk more about acids in Ch 19)
There are two types of acids:
1. binary acid: contains hydrogen and one other
element
-when naming use the prefix hydro- plus the root of
the second element with the suffix –ic, followed by
the word acid.
-ex: HCl
H = hydroCl = chloride = chloric
hydrochloric acid
Some acids are not binary, but are named according to
the binary acid rules when oxygen is not present, as in
HCN.
H = hydro
CN = cyanide = cyanic
hydrocyanic acid
2. oxyacid: an acid that contains an oxyanion (oxygen
containing polyatomic ion)
-the name depends on the oxyanion present
-the name consists of the root of the anion, a suffix,
and the word acid
♦if the anion suffix is –ate, it is replaced with -ic
♦if the anion suffix is –ite, it is replaced with -ous
-examples:
~HNO3
NO3 = nitrate
= nitric
nitric acid
~HNO2
NO2 = nitrite
= nitrous
nitrous acid
Name the following acids:
1. HBr
2. H3PO4
3. H2SO4
4. H2SO3
5. H2CO3
Use the prefixes in the molecule’s name to determine the
subscript for each atom in the compound.
- phosphorus tribromide
P
Br
1 (no prefix)
3 (tri)
PBr3
- the formula for an acid can be derived from the
name as well
♦charge of the oxyanion or anion gives the number
of hydrogens
hydrofluoric acid = HF
(1 H because fluorine has a -1 charge)
1. oxygen difluoride
2. dinitrogen tetrasulfide
3. phosphorus pentachloride
4. iodic acid
5. phosphoric acid
TEST # 2
Metallic bonds are similar to ionic bonds because they
often form lattices in the solid state.
-eight to twelve metal atoms surround another, central
metal atom
Instead of sharing electrons or losing electrons, the outer
orbitals overlap.
-electron sea model: all metal atoms in a metallic
solid contribute their valence electrons to form a ‘sea’
of electrons around the metal atoms.
-valence electrons are free to move from atom to
atom (delocalized electrons), forming metallic
cations
metallic bond: attraction of a metallic cation for the
delocalized electrons that surround it
This bonding contributes to the unique properties of
metals:
1. generally have high melting and boiling points, with
especially high boiling points
-due to the amount of energy needed to separate
the electrons from the group of cations
2. malleable (hammered into sheets) and 3. ductile
(drawn into wire)
-mobile electrons can easily be pulled and pushed
past each other
4. durable
-though electrons move freely, they are strongly
attracted to the metal cations and are not easily
removed from the metal
5. good conductors
-free movement of the delocalized electrons,
allowing heat and electricity to move from one
place to another very quickly
6. luster
-interaction between light and delocalized
electrons
As the number of delocalized electrons increases, as in
transition metals (d electrons), the hardness and
strength also increases.
-alkali and alkaline earth metals are soft (s valence
electrons only)
It is easy to combine 2 or more different metals to make a
metallic crystal
-alloy: mixture of elements with metallic properties
-the properties of alloys differ from those of the
individual elements that make it up
1.
2.
3.
4.
5.
6.
7.
8.
9.
What is a metallic bond?
What is an alloy?
Describe the electron sea model.
What occurs with orbitals in metals?
How is metallic bonding similar to ionic bonding?
What are delocalized electrons?
What contributes to a metal’s high boiling point,
malleability, ductility and conductivity?
List the other 2 properties of metals.
What happens to strength and hardness as you
decrease the number of delocalized electrons?