Chapter 14: Liquids and Solids (or States of Matter)
Goals:
 Learn about Van der Waal’s forces (London-dispersion forces, dipole-dipole attraction, and
hydrogen bonding), and understand the effects these intermolecular forces have on solids
and liquids.
 Explore water – including how hydrogen bonding accounts for many of its properties.
 Understand what is occurring during phase changes, and how to use the heat of
vaporization, heat of fusion, and heat capacities of a substance to perform calculations
about the energy changes that occur with phase changes.
 Understand the relationship between vaporization (both evaporation and boiling) and vapor
pressure.
 Relate the boiling point of substances (including water) to their vapor pressure.
 Learn about the various types of solids, and how these types result from different types of
bonding and intermolecular attractions.
As you’re no doubt aware, the three states of matter we work with in chemistry are solids, liquids,
and gases. In the next few classes we’ll explore how these states work on a microscopic level.
Table: General properties of the three states of matter
property
solid
liquid
gas
density
high
high
Very low
compressibility
low
low
Very high
volume
constant
constant
fills container
shape
constant
Takes shape of
Takes shape of
container
container
Somewhat random / completely random /
structure
Neat, orderly
disordered
disordered
We’ll discuss why each state has each of the properties above.
Gases (a quick review)
The Kinetic Molecular Theory of Gases:


Definition: The kinetic molecular theory (KMT) of gases says that the properties of gases are
determined by the interactions of the gas atoms/molecules with each other.
Basic postulates of the kinetic molecular theory:
o
Gas particles are infinitely small:
 In a gas, the gas molecules are separated from one another by a lot of empty space.
As a result, the gas particles can be said to have negligible volume compared to the
overall volume of the gas. This has the advantage of making the math much easier.
o
Gas particles are in constant, random motion:
 In a gas, the particles constantly move very quickly all over the place, changing
direction only when they hit something and bounce off. Energy is transferred between
particles when they collide, but this energy is conserved (not lost).

The collisions of gas particles with the sides of the container they’re in is called
pressure.
1
o Gas particles don’t experience intermolecular forces:
 Because the gas particles are so small, are so far apart, and are moving so quickly, the
molecules don’t interact with each other much. To make our lives (and the math
easier), we just ignore the very little interaction that does take place.
o The kinetic energies of gas particles are proportional to their temperatures (in Kelvin).
 This should make sense: All this says is that the hotter the particles in a gas are (the
faster the average speed of the particles), the more energy the gas particles have.

Why Kelvin? Because if you used degrees Celsius, you’d get negative energy
whenever the temperature dropped below the freezing point of water (0 o C). As a
result, we have to use a temperature scale where the zero point energy of a molecule
corresponds to zero on the temperature scale.

This relationship is mathematically described by: (this formula should look familiar)
1
KE  mv 2
2
Properties of Gases:

Gases have low density: This makes sense: If there’s lots of space between gas
molecules, then you’d expect low density.

Gases can be compressed and expanded:
o You can compress a gas because all you’re doing is just squishing the particles
together.
o Gases expand because the particles move all over the place unless you stop them
from doing so.

Gases diffuse: Diffusion is when two things put in the same container mix together.
o For example, if you open a jar of pickles in a room, you can soon smell pickles all over
the place – this is because the pickle smell diffuses throughout the room.

Gases effuse: Effusion is when a gas escapes through a hole in a container into a vacuum.
o Because of ½ mv2, light molecules (those with a low molar mass) move more quickly
than heavy ones (if they are at the same temperature). As a result, light molecules
diffuse and effuse more quickly than heavy ones.
o Graham’s Law of effusion:
Rate A

Rate B
MW A
MW B
o Sample problem: If the scent molecules in pickle juice have a molar mass of 450 g/mol
and the scent molecules in ammonia have a molar mass of 17 g/mol, how much faster
will ammonia molecules move than pickle molecules?
5.15 (21.21/4.12)
2
Liquids: The Joy of Intermolecular Forces
Whereas solids are materials in which the particles are very tightly stuck to one another, liquids
are materials in which the particles have a little more interaction.
As you already know from experience, liquids are materials that can easily change shape but also
tend to have fairly high densities (nearly the same as solids). Why do the molecules hang
together at all? The answer: Intermolecular forces.
Intermolecular forces - Interactions that hold the particles in a liquid together. There are
three types of intermolecular forces (often called Van der Waals’ forces) that we need to
consider. (All three are varying degrees of attraction between partial positive charges and
partial negative):

Dipole-dipole forces: Interactions in which polar molecules stick to each other like little
magnets when the partial positive side of one is attracted to the partial negative side of
another.
o The more polar the molecule, the stronger the attraction!

Hydrogen bonds: A very strong dipole-dipole force that occurs when the lone pair
electrons on O, F, or N interacts strongly with a hydrogen atom bonded to O, F, or N.
o Essentially, these bonds are so polar that the lone pair electrons (negative) on one
molecule want to stick to the very positive hydrogen atoms on another molecule.
o The more hydrogen bonding that a molecule can do, the stronger this force is.

Water has a MP of 00 C, while methanol has a MP of -980 C.

London Dispersion Forces: When nonpolar molecules are attracted to one another via
temporarily induced dipoles.
o Essentially, nonpolar molecules stick together magnetically – like polar molecules.
How can this work?
o The bigger the molecules, the stronger the force (because there are more electrons
to become unbalanced and interact with each other).
3
o Why are intermolecular forces important?

The stronger the intermolecular force, the higher the melting and boiling points.
o Because melting and boiling both involve the movement of particles from their
neighbors, anything that causes neighboring particles to stick together will raise them.

Intermolecular forces make liquids almost as dense as solids:
o Because the intermolecular forces in liquids keep the molecules stuck to each other,
liquids are nearly as dense (and in a few cases even denser) than solids.

Intermolecular forces allow liquids to flow – a property called “fluidity”.
o Because the molecules in a liquid are attracted to each other but not permanently stuck
in place, the molecules can move from one place to another – flowing!

Intermolecular forces cause differences in the viscosities of liquids:
o Definition: Viscosity is the resistance to a liquid flowing. High viscosity = slow flowing.
o The higher the intermolecular force, the more viscous the liquid. This is because
molecules that are held tightly together want to move apart less than molecules that are
loosely held.
o Viscosity also goes down with increasing temperature (i.e. things flow more easily at
high temperatures) – this is because the energies of the particles in the liquid are
beginning to get to the point where they can overcome some of the attractive forces.
Intermolecular forces cause surface tension in a liquid.
o Definition: Surface tension is the energy needed to increase the surface area of a
liquid – the higher the surface tension, the harder it is for something to push through
the surface of a liquid.
o Stronger intermolecular forces cause the surface molecules to hold together more
tightly, making the surface tension higher.
o This is why some things that are heavier than water (i.e. water bugs, leaves, etc) don’t
fall through.

Intermolecular forces cause capillary action.
o Capillary action: The tendency of some liquids to rise when placed in a small tube –
this explains why putting one edge of a paper towel will eventually cause the whole
towel to get wet, as well as how water gets to the top of a tall tree.
o This happens because water molecules want to grab the surface of the walls of the
tube with their intermolecular forces more than they want to grab each other. This
causes them to move up the sides of the tube (away from each other).
o This causes the meniscus.
o Eventually, the pressure of the water height overcomes the attraction of the water for
the sides of the tube and the water stops rising.
4

Ranking of intermolecular (and intramolecular) forces (strongest to weakest):
o Ionic interactions and covalent bonds (they are not intermolecular forces, they are
intramolecular forces) – it takes a huge amount of energy to break these.

Example: The BP of NaCl is 14130 C.
o Hydrogen bonding

Example: The BP of water is 1000 C.
o Dipole-dipole forces

Example: The BP of H2S is -59.60 C.
o London dispersion forces

Example: The BP of methane is -161.50 C.
Questions that may be asked:
 What intermolecular force is molecule X likely to experience?

o
Draw the Lewis structure of the molecule.
o
If the molecule is nonpolar, it’s Van der Waals forces.
o
If the molecule has H bonded to O, N, or F, it’s hydrogen bonding.
o
If the molecule is polar but H isn’t bonded to O, N, F, it’s dipole-dipole forces.
o
[do some examples: CH4, CH4O, HCN]
Rank the following by increasing melting/boiling point.
o To solve, determine the type of intermolecular force that each molecule is undergoing.
o You may assume that molecules with stronger intermolecular forces will have higher
MP and BP than those with weaker ones.
o
Example: Rank CH2O, CO2, and CH2O2
(ranking intermolecular forces worksheet)
5
Phase Changes:
Solid  Liquid transformations:




Melting is when a solid becomes a liquid. The reverse of this process is called freezing.
Melting and freezing are the reverse of one another and happen at the same temperature.
Why things melt:
o Solids melt when the amount of energy that’s available (because we’ve heated them) is
greater than the amount of energy that’s holding them together.
 Covalent compounds melt at low temperatures because the amount of energy that
holds the particles together through intermolecular forces is very small.
 Ionic compounds melt at high temperatures because the lattice energy that holds
the ions together is very high.
Why things freeze:
o Liquids freeze when enough energy has been taken away from the liquid that the particles
are no longer able to stay separate – the intermolecular forces (or lattice energy, in the
case of ionic compounds) now force them to combine into a solid.
Liquid  Gas transformations: Vaporization and Condensation


In a liquid, the particles move into the gas phase when they get enough energy to break
free of the intermolecular forces that hold them together.
Vaporization – liquid to gas (can be evaporation or boiling)
o Evaporation is the process in which only a very few of the molecules in a liquid have
enough energy to break free. It occurs below the boiling point, and only at the surface.
 The pressure of the molecules that have become a gas is called the vapor
pressure of the liquid.
 The higher the temperature of the liquid, the higher the vapor pressure (because
more of the molecules have gotten enough energy to become a gas)
o Boiling: When the vapor pressure of the liquid becomes equal to the atmospheric
pressure, the molecules in a liquid have gotten enough energy to break free of the
intermolecular forces that hold them together.
 Boiling point: The temperature at which this happens (varies with pressure).

Condensation: When enough energy has been removed from a gas that intermolecular
forces again hold them together as a liquid.
o Condensation and vaporization are the reverse process of one another and happen
at the same temperature.
6
Solid  Gas transformations: sublimation and deposition


Sublimation is when things go directly from the solid phase to the gas phase. This happens
with dry ice (when the white gas comes off of the block).
o This is why ice cubes in the freezer get smaller over time.
o This is what causes “freezer burn.”
o This is how things are freeze-dried (though at pressures below 1 atm)
Deposition is when things go from the gas phase directly to the solid phase.
o This is why frost sometimes builds up on the sides of a freezer.
o Snow and frost are deposition.
Phase diagrams: How we figure out what’s going on

Phase diagrams: Diagram that shows you what phase changes occur at different
temperatures and pressure.
o Obviously, phase changes take place when the temperature changes – that’s how we
normally boil water and melt ice.
o However, pressure is also important – remember how things boil if the vapor pressure
= the atmospheric pressure? Well, if we decrease the pressure inside a container,
things boil at lower temperatures.

Pressure changes have a similar effect on other phase changes.
Important features of phase
diagrams:




Lines: Along the lines that
separate the phases, both phases
are equal to stably coexist. That’s
why you can put a glass of ice
water in the refrigerator and find
both the ice and liquid water there
after a few days.

Normal freezing point: The
temperature at which a substance
freezes/melts at a pressure of 1
atm.
Normal boiling point: The temperature at which a substance boils/condenses at a
pressure of 1 atm.
Triple point: The conditions of pressure and temperature at which all three phases of
matter can stably coexist. For water this is 0.006 atm and 0.01 0 C, which makes it
impossible to observe without special equipment.
Critical point: The temperature above which water cannot exist as a liquid. Above this
temperature water exists as something between a liquid and a gas.
7
Some terms and ideas you need for studying thermodynamics:
As with everything, thermodynamics has specialized terms to describe the things that go on in the
real world. Before moving on to seeing how energy behaves, we need to first understand the
terms that are used to describe it.
Energy: The ability to do work or to produce heat. There are two types of energy:

Kinetic energy: The energy something has when it moves. (i.e. moving objects, moving
particles, vibrating molecules, etc.)
o Temperature is a measure of the particles in an object. We know this from the
KMT, which says that the amount of energy is proportional to the temperature (in K).
The more the particles in an object move around, the higher the temperature.

Potential energy: Stored energy that’s waiting for its chance to get moving. (i.e. objects
that are waiting to fall off of a shelf, energy stored in chemical bonds, etc.).
o Chemical potential energy: The energy that’s stored in chemical bonds.
o Heat (q): The movement of energy from one thing to another through the motion of molecules
(thermal energy).

Heat spontaneously moves from hot things to cold. This is why a hot pan can burn you
and you can’t burn a hot pan – the energy goes only from the pan to you because it’s
hotter.

Heat and temperature are NOT the same thing: Heat is the transfer of energy,
temperature is a measure of the kinetic energy of the object once the energy has
finished transferring.
Quantifying energy:

The traditional unit of energy is the calorie (cal), which is the amount of energy you need
to add to 1 gram of water to heat it by 10 C.

Food is measured in units of 1000 calories called kilocalories (kcal), which is more
commonly known as the Calorie (Cal).

The metric unit of energy is the joule (J). There are 4.184 J/cal.

Because a joule isn’t very much energy, we usually measure energy in units of 1000 joules
called kilojoules (kJ).
(In class practice worksheet about energy, followed by more notes)
8
Enthalpy: ∆H
Now that we’ve learned what heat is, it’s handy to think about how much heat a system can
potentially give off to other systems. This term is called “enthalpy.”
Enthalpy (H): The amount of heat that a system can potentially give to other systems.
 Unfortunately, it’s impossible to know what the overall enthalpy of a system is – after all, how
much heat one system can give to another depends on a large number of factors, including
the nature of the system and the nature of whatever system it wants to give energy to.
o Instead of talking about how much enthalpy something has, we instead talk about how
much the enthalpy of a system changes when heat is taken away from it or added to it.
o This term is given the symbol ∆H (“delta H”), where ∆ represents the change in
enthalpy that occurs during a process.

Enthalpy changes having to do with phase changes:
o Heat of vaporization (∆Hvap) is the amount of heat required to convert 1 gram (or 1 mole)
of a liquid at its boiling point into vapor without an increase in temperature. Watch the
units! (may be called the Molar heat of vaporization if it is the heat required for 1 mole.)
o Heat of fusion (∆Hfus) is the amount of heat required to convert 1 gram (or 1 mole) of a
solid at its melting point into a liquid without an increase in temperature. Watch the units!
(may be called the Molar heat of fusion if it is the heat required for 1 mole.)
o Heat capacity is the heat required to raise 1 gram (or 1 mole) of substance by one degree
Celsius (or one Kelvin). (Molar heat capacity is per one mole; while Specific heat is per
1 gram.) Heat capacity is often shown as C; and often with a subscript p indicating that the
pressure is held constant.
o The heat capacity of a substance varies depending on which phase it is in. The specific
heat of liquid water is 1 calorie/gram °C = 4.184 joule/gram °C Watch the units!
Enthalpy changes caused by heating/cooling:
How much energy (heat) is needed to heat something up (or cool it down)?
The enthalpy change that accompanies the heating/cooling of a pure substance is determined by
the equation:
∆H = mCp∆T
where:
o ∆H = the change in enthalpy (positive for heating, negative for cooling)
o m = the mass of the thing being heated (in grams)
o Cp = the specific heat / heat capacity – the amount of energy needed to heat the
substance by 10 C.
o ∆T = the change in temperature (in degrees Celsius).
9
Sample problems:
o How much energy will be needed to heat 45 grams of ethanol (Cp = 2.44 J/g0C) from 200 C
to 400 C? 2196 J
o If burning a single match gives off 250 J, how much hotter can it make a 50 gram block of
silver (Cp of silver is 0.235 J/g0C) 21.30 C
(Lab: Calculating the heat capacity of water, or calculating the heat of fusion of water.)
(Heating/cooling homework)
Enthalpy changes during phase changes:
As you might imagine, phase changes also change the enthalpy of a system. As you know,
melting a compound causes the energy of the particles to increase as they break intermolecular
forces partially, and boiling breaks them completely.
Freezing/melting:
The amount of energy that’s added/removed from a substance during the freezing or melting
process is described by the equation:
∆H = n ∆Hfus
where
o ∆H = the enthalpy change for this process.
o n = the number of moles of the compound melting or freezing
o ∆Hfus = the molar heat of fusion (which is a constant for a given pure substance)
Example: How much energy is required to melt 56 grams of frozen water at 0oC? (∆Hfus water =
6.01 kJ/mol) 18,700 J or 18.7 kJ
Important: When undergoing the phase transition from liquid to solid (or vice versa), all of the
energy goes into breaking intermolecular forces. As a result, the temperature of the material
doesn’t change as it undergoes the transition! (You saw this in your lab before break.)
o This is why ice cubes keep a cold drink at 00 C until they have completely disappeared.
10
Boiling/condensing:
The amount of energy that’s added/removed from a substance during the boiling or condensing
process is described by the equation:
∆H = n ∆Hvap
where
o ∆H = the enthalpy change for this process.
o n = the number of moles of the compound boiling or condensing
o ∆Hvap = the molar heat of vaporization (which is a constant)
Example: How much energy is required to boil 56 grams of frozen water? (∆Hvap of water = 40.7
kJ/mol)
126.6 kJ = 126,600 J
Important: When undergoing the phase transition from liquid to gas (or vice versa), all of the
energy goes into breaking intermolecular forces. As a result, the temperature of the material
doesn’t change as it undergoes the transition!
o This is why boiling water remains at 1000 C until all of the water has been boiled away
rather than increasing in temperature!
Now that we’ve learned all of this, we have enough information to learn what the energy
change is for a compound as we heat it from a frozen material through both phase
changes until it’s a gas.
o The whole process of figuring out how much energy is required to make a chemical or
physical change take place is called calorimetry. The device we use to do this is called a
calorimeter.
o To do this, you must first figure out where you are on this graph:
11
Once you know where you are and where you need to end up, you can figure out what equations
you need to figure out the enthalpy change for the process.
Example: How much energy is needed to heat 55.0 grams of water ice from a temperature of
-15.0 0 C to steam at a temperature of 150.0 C?
Solution: This problem requires several steps:
o Step 1: Heat the ice from -150 C to 00 C (Cp of ice = 2.03 J/g0C) (Watch the units!!!)
o ∆H = mCp∆T =
(1,675 J)
o Step 2: Undergo the phase change from a solid to liquid. (∆Hfus of water = 6.01 kJ/mol)
(Watch the units!!!)
o ∆H = n∆Hfus =
(18.39 kJ)
o Step 3: Heat the liquid from 00 C to 1000 C Cp for liquid water = 4.184 J/g0C
o ∆H = mCp∆T =
(23,012 J)
o Step 4: Undergo the phase change from a liquid to gas. (∆Hvap of water = 40.7 kJ/mol)
o ∆H = n∆Hvap =
(124.54 kJ)
o Step 5: Heat the steam from 1000 C to 1500 C (Cp of steam = 2.01 J/g0C)
o ∆H = mCp∆T =
(5,528 J)
And when you’re done with all of this, just add up the different values you found above:
∆Htotal = 1.68 kJ + 18.39 kJ + 23.01 kJ + 124.54 kJ + 5.53 kJ
∆Htotal = 173.15 kJ = 173 kJ
(3 significant figures)
12
Solids
The single characteristic that defines the properties of solids is this: In a solid, the particles don’t
move around very much. This is because, for various reasons, the particles in a solid are locked
into place – so all they can do is vibrate around a fixed position). As a result, solids are generally
quite dense and hard.
The structure of solids: There are two main types of structure that solids have (crystalline
and amorphous):
 Crystalline solids: Solids in which the molecules/atoms are stacked together in orderly and
repeating 3-D structures.

Each of the repeating units that a solid is made up of is called a “unit cell” – these unit cells,
when repeated over and over again, give the structure of the entire crystal.

The three types of crystalline solid: There are 3 main types of crystalline solids:
Ionic, Molecular, and atomic. Within the Atomic solids, there are 3 subtypes: Network,
Metallic, Group 8A

The type of solid is determined by the intra/intermolecular forces holding the substance
together.

Ionic solids: Solids in which ions are held together via electrostatic attractions (positive
and negative ions held together like strong magnets). Examples: NaCl, Li2CO3, etc.
o Properties: (see pages 14-15).

Molecular solids: Solids in which covalent molecules are locked into crystals by various
intermolecular forces (London dispersion forces, dipole-dipole forces, or hydrogen
bonds). Examples: water, ammonia, carbon dioxide, Sulfur (S8), white phosphorous
(P4), etc.
o Properties: Low/medium MP/BP, soft. (see page 16)

Atomic solids
 Covalent network solids: Solids in which atoms bond together covalently to form
crystals – these are basically great big molecules. The bonds between molecules
are the same as the bonds within them. You do not find individual SiO2 molecules
within a quartz crystal. Examples: diamond (C), quartz (SiO2), etc.
o Properties: High MP/BP, hard, brittle. Often indistinguishable from ionic
compounds; except that they almost never conduct electricity in any form.

Metallic solids: Solids in which delocalized mobile electrons hold the metal atoms
together. Examples: Any metals.
o Properties: Usually high MP/BP, malleable/ductile, conductive, shiny, and hard –
though these can vary greatly depending on the strength of the metallic bond.
(page 17)

Group 8A (Noble gases):
o At extremely low temperatures (and high pressure), noble gases will become
solids held together by weak London dispersion forces.
13

Amorphous solids: Solids that have no regular structure (i.e. they don’t form crystals).







In these solids particles are randomly arranged in three dimensions.
They don’t have sharp melting points.
Amorphous solids are formed due to sudden cooling of liquid.
Amorphous solids melt over a wide range of temperature
Examples:
This frequently occurs when liquids solidify extremely quickly - because it takes time for
the particles in a crystal to orient themselves properly, a very rapid cooling may cause this
crystallization process not to work.
 The resulting solids are typically hard and very brittle. Example: Glass.
It sometimes also happens (especially in organic compounds) that very long molecules
tangle up with each other to form very soft and pliable solids.
 Examples: Rubber, plastics
Notes: Properties of Ionic Compounds
1)
Salts have ordered packing arrangements.

The anions and cations stack in regular arrangements to minimize the distances between
the anions and cations.

The three dimensional pattern in which they end up stacking is called a “crystal lattice” or
“crystal.”
There are many different types of
lattice. This is the simplest type,
referred to as “simple cubic.”
2)

Salts conduct electricity when dissolved or melted.
Liquids that conduct electricity are called electrolytes.

All particles in an ionic compound have either positive or negative charges – these charges
make it possible for electricity to be conducted in molten salts or salt solutions.

Electricity is not ever conducted when salts are in the solid state, because the ions are
locked into their lattices.

If a salt doesn’t dissolve in water (and many don’t), it won’t conduct electricity even if it’s
sitting in water. It needs to dissolve for the ions to move around freely.
14
3)
Salts have a high melting and boiling point.

When ionic compounds melt, the ions are heated until there’s so much energy that the
forces holding them together break apart and the ions move freely around as a liquid.

Because there’s so much energy holding ions next to each other (called the lattice
energy), you need a lot of heat (and very high temperatures) to make this happen.

Melting points range from under 1000 C to over 10000 C (depending on how tightly the ions
are stuck to each other). Generally, it takes a few minutes for a Bunsen burner to melt a
salt (if it melts at all!), though there are exceptions.
4)
Salts are hard and brittle.

Salts are very hard because the ions are stuck together in a rigid lattice. All ions hold
neighboring ions in place very tightly – as a result, a small applied force won’t move the
ions, causing the compound to feel hard.

Salts are brittle because once you put enough energy into the lattice to make it come
apart, you’ve put enough energy into it to destabilize the whole lattice. This causes it to
shatter (or at least crack over a large area).
5)
Ionic compounds rarely burn.

6)
Compounds need to contain carbon and hydrogen before they burn. Because both are
only rarely found in ionic compounds, most ionic compounds don’t burn.
The formation of ionic compounds is always exothermic (gives off heat).

As mentioned earlier, elements go to lower energy states when they become like their
nearest noble gases.

All of the extra energy they lost has to go somewhere – out in the form of heat. As a
result, very large amounts of heat are frequently given off when ionic compounds form.
o Demo: Burn magnesium ribbon

This stability is reflected in the lattice energy of the compound – the amount of energy
needed to pull apart the ions in one mole of an ionic compound.
o The smaller the ions, the more negative (and more stable) the lattice energy,
because the charges are closer to each other.
o The bigger the charges on the ions, the more negative (and more stable) the lattice
energy, because the bigger charges cause the ions to stick together more tightly.
15
Properties of covalent compounds / Molecular solids:
All properties of covalent compounds are determined by the fact that covalent compounds form
molecules, while ionic compounds form crystals.
 A good model for thinking of this:
o Ionic compounds are like strong magnets all locked together into a big block.
o Covalent compounds are like rubber balls thrown together into the same bucket.
1)


Covalent compounds have low melting and boiling points:
In ionic compounds, all of the atoms are magnetically stuck together in great big crystals.
-
-
+
To melt an ionic compound, you
need to overcome the magnetic force
for all the ions in the crystal.
In covalent compounds, the molecules only have very weak forces (called Van der Waals
forces) holding them to each other. As a result, covalent compounds can be found as
solids, liquids, or gases at standard room conditions.
δδ+
δ-
2)
+
δ+
δ-
δ+
To melt a covalent compound, you DO NOT
BREAK COVALENT BONDS!!!! All you do is
separate the molecules from each other. Since the
Van der Waals forces holding them next to each
other are very weak, it doesn’t require much
energy to do this.

Covalent compounds are soft and squishy
This is only a general property – some covalent compounds can be quite hard.

Covalent compounds usually don’t dissolve in water as well as ionic compounds.
Explain how water pulls salts apart and show briefly how the water molecules have less
success grabbing onto covalent compounds.

Covalent compounds don’t conduct electricity (neither when solids, melted, nor
when dissolved in water)
Since there are no charged particles (as in ionic compounds) and no delocalized electrons
(as in metals), they don’t conduct electricity at all.
3)
4)
5)


Covalent compounds sometimes burn.
Compounds that burn usually contain carbon and hydrogen. Because carbon and
hydrogen form covalent molecules when they bond with each other, some covalent
compounds are able to burn.
Covalent compounds that don’t contain carbon and hydrogen usually don’t burn.
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Metallic bonds and properties
How metals bond:
Electron sea theory: Metals consist of positively charged “cations” (thought they are not truly
cations) floating around in an ocean of constantly moving electrons (“the sea”). The electrons in
the sea are the outer electrons (not technically the valence electrons – see property #6 on the
next page.)
 Unlike ionic compounds, the electrons don’t get transferred from one atom to another.
Instead, they move between all of the different positively charged nuclei in the metal.
 Because the electrons move around, “mobile”, they are referred to as being delocalized
electrons.
Metallic bonds are the forces that hold positively charged metal cations within this ocean of
delocalized (mobile) electrons.
How this affects the properties of metals:
1)

2)
Metals form “mobile lattices”:
Because the cations are held together by a bunch of electrons that move around, the
cations can also move around when (when pounded, pulled, etc.) – and the electrons will
follow along as needed.
Metals have high boiling points (but differing melting points):

To melt a metal, you just have to make it so that the cations can move around each other.
Depending on how many electrons you have, this can either be very easy or very difficult:
o Hg – which is actually Hg2 molecules, is a metal that does not exhibit true metallic
bonding – hence its unusually low melting point
o If the electrons don’t hold the metal nuclei together very tightly in place, the metal
has a low melting point. Lead and some alloys have relatively low melting points.
o If the electrons do hold the metal nuclei together tightly, then the metal has a high
melting point (Fe).

To boil a metal, you not only have to move the cations, but have to separate them entirely.
Because the electrons bind with the cations, it’s very hard to get them all to move away
from each other.
o So most metals have high boiling points, even if they have low melting points.
3)

Metals are malleable and ductile:
Malleable: They can be pounded into sheets (they are bendy).

Ductile: They can be pulled into wires.

Why? As you move the metal cations around, and the electrons follow along as needed.
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4)


5)
Metals are good conductors of electricity and heat:
Since moving electrons = electricity, and metals have lots of delocalized (mobile)
electrons, metals easily conduct electricity.
Since heat can also be transmitted easily when electrons are mobile (because lots of
collisions occur, and energy is conducted through collisions), metals are also good
conductors of heat.

Metals are shiny:
This has to do with how light messes with moving electrons. It’s also too complicated to
talk about here, so just know that delocalized electrons are responsible for this one, too.

Transition metals are hard:
The more delocalized electrons you have, the harder the metal.
6)

Because transition metals have both s- and d-electrons that can move around, they are
harder than alkali (group 1) and alkaline earth (group 2) metals that have only s- electrons
available to be delocalized.

Honors only: Metals in the p-block (post-transition metals) are fairly soft as well because
the p-electrons are the only ones that really do much moving around.
7)



Metals form alloys:
Alloys are metallic mixtures which contain at least two elements – at least one of which is a
metal.
These elements are chosen to give metals desirable properties (increased strength, higher
MP or BP, etc). This is why steel (which usually contains C, Cr, Ni, or other elements
alloyed in iron), is preferred over pure iron.
Types of alloys:
o Substitutional alloy: Alloys in which one of the nuclei is replaced by another
element’s nucleus of similar size. (18 carat gold, brass):
o The additional elements usually provide more delocalized electrons to the sea.
o Interstitial alloy: Alloys in which a much smaller nucleus fits into the spaces
between the existing nuclei (steel):
o The extra elements form strong directional bonds with some of the metal cations.

Some alloys (e.g. stainless steel) are both substitutional and interstitial alloys.
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