Transcript Document

Chapter 1
AP Chemistry
Milam
1-1 matter and energy
• Matter takes up space and has mass, so
anything besides light, energy, forces
• Mass is how much matter something
consists of
• Energy is the ability to do work (useless
definition)
• Kinetic energy ½ mass * velocity2
1-1
• Potential energy is stored energy and there
are multiple kinds of potential energy
(nuclear, gravitational, chemical…)
• Exothermic is a process that releases heat
(surroundings get warmer)
• Endothermic is a process that absorbs heat
(surroundings get cooler)
1-1
• Laws of conservation – In most situations,
both energy and mass are conserved,
meaning if you start with a certain amount
of energy or mass you end up with that
much
• However, mostly in nuclear reactions,
energy and mass may interchange according
to the equation Energy = mass * (speed of
light)2
• E= mc2
1-2 Chemistry – A Molecular
View of Matter
• Chemistry studies very small objects (atoms
and molecules), but can do so on a large
scale and then relate this back to a small
scale
• Democritus was one of the first people to
postulate the idea of atoms, although he
performed no experiments
• John Dalton did perform experiments….
1-2
• Dalton discovered that:
– Every element is made up of atoms
– All atoms of an element have identical
properties
– Atoms are conserved
– Compounds form in small whole-number ratios
– The relative # and type of atoms are constant in
a given compound
1-2 about the atom
• Atoms are made of electrons, protons and
neutrons
• Neutrons are slightly bigger than protons
and both are much larger than electrons
• Neutrons and protons make up the nucleus
and electrons are in the electron cloud
which takes up the majority of the volume
of an atom
1-2
Atomic #
• Atomic # (Z) is the number of protons in an
atom.
• Because the atoms of an element all have
the same # of protons they also have the
same atomic #
• You can locate an element on the periodic
table using atomic # (ex. Hydrogen Z=1)
1-2
Molecules
• A molecule typically consists of more than
1 atom bonded together that is stable at
normal conditions and electrically neutral.
• There are seven diatomic molecules, these
exist in nature rather than the elemental
forms (H2, N2, O2, F2, Cl2, Br2, I2)
1-3 states of matter
• Solids, liquids and gases will all be
developed in more detail in the future, for
now be able to recognize each by the
pictures at the molecular level
1-4 Chemical and Physical
Properties
• Chemical properties are how a substance
will and will not react with other chemicals
– A chemical property of sodium is that it reacts
violently with water
• Physical Properties are any other property
– Color, density, hardness, melting point, shape,
solubility…..
– Ice melts at 0° C is a physical property of water
1-4
• Extensive properties depend on the amount
of the material
– Volume, mass, weight….
• Intensive properties are independent of
amount
– Color, melting point, density, chemical
properties….
1-5 chemical and physical
changes
• Chemical changes involve a chemical
reaction
• Signs of a chemical change include change
in color, gases or light being produced,
energy being absorbed and released,
forming a precipitate
• Physical changes involve any change where
the chemicals remain the same
1-5
• Examples of physical changes include:
phase changes (melt, sublimate, condense),
dissolving, breaking, ripping….
• Physical changes would not protect you
from poison, chemical changes might or
might not depending on the new product
1-6 Mixtures, substances,
compounds and elements
• Most everyday items are mixtures, there are
two kinds of mixtures
– Heterogeneous – where the mixture is not
uniform
– Ex. Salad, Legos, sand
– Homogeneous (aka solution) – where the
mixture is uniform
– Ex. Kool-aid, alloys, air
1-6
• An element is a type of atom with a specific
number of protons, an element cannot be
decomposed into a simpler substance
– Ex. Br2, Cu, He, Silver metal
• A compound is a substance with two or
more types of atoms in a fixed ratio of
amounts
– Ex. Water H2O, C6H12O6, H2SO4
1-6
• The law of definite proportions (law of
constant composition) says that if you take a
compound and compare the percent mass of
each element, they will always be the same
for each compound
• So if you analyze pure water to be 88.9%
oxygen and 11.1% hydrogen, it will always
have those percentages.
1-6
• The table with elements and symbols on
page 18 is worth becoming familiar with
• Spend 5 minutes 2-3 times looking over the
list and memorizing the elements and
symbols that you aren’t already familiar
with
1-7 Measurements in Chemistry
• Being familiar with some of the prefixes in
table 1-6 on page 19 is important,
specifically
– Kilo, Centi, milli, micro and nano are
commonly used
– Pico, mega, deci are used in some odd
situations, but when those arise you can just
look them up
1-8 Units of Measurement
• You are probably already familiar with most
of the units, but maybe you are not used to
using some of the SI units
• You need to know: length – meters; Mass –
kilograms; Volume – liters; time – seconds;
temperature – kelvins; amount of substance
- mole
1-9 Use of Numbers
• Scientific notation is a way of writing really
small and really large numbers
• The basis for it is that 108 = 100,000,000 so
if I am doing a calculation and get an
answer of 200,000,000 I can write 2x108
instead
• Be able to convert between numerical and
scientific notation
1-9 sci. not.
• Ex. 140,678 into scientific notation
• 1.40678 x 105
• Move the decimal until the # is between 1
and 9.9999……
• Count the number of decimal places moved
and make that the exponent
• Ex. 2, turn 4.57 x 10-7 into numerical
notation
1-9 sci. not.
• Ex. 2, turn 4.57 x 10-7 into numerical
notation
• 0.000 000457
• Negative exponents give numbers less than
1 and positive exponents give numbers
greater than 1
• 100 = 1
1-9 sig figs
• Significant figures tell you how accurate of
a measurement was made
• To take a reading, you read the smallest
increment on the measuring tool and
estimate one additional place
• Once you take this reading you need to keep
the amount of significant figures consistent
when you make calculations from your data
1-9 sig figs
• In this class we assume that all calculations
are from lab data, hence, all calculations
must keep significant figures intact
• Significant figures are worth points on the
free response sections of the test and could
show up on the multiple choice as well
• All digits 1-9 are always significant, 0’s
depend on the situation
1-9 sig figs
• 0’s that are in between two non zero digits
are always significant (ex. 303 is 3 sig figs)
• If there is no decimal:
– All 0’s are not significant
• If there is a decimal:
– All 0’s to the right of the last nonzero digit are
significant
1-9 sig figs
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Ex. 1 - 05000 has 1 significant digit
Ex. 2 – 6470.0 has 5 sig figs
Ex. 3 – 0.004 has 1 sig fig
Ex. 4 – 0.00320 has 3 sig figs
Ex. 5 – 0.0040300 has 5 sig figs
Sometimes it helps to read from right to left
with a decimal, and the last nonzero number
is the last sig fig
1-9 sig figs
• In addition and subtraction the significant
figures are determined by the least
significant place
• Adding 4 + 0.78 gives 4.78, but since 4 only
has a significant figure in the ones place, the
answer needs to be good only to the ones
place as well
• 4 + 0.78 = 5
1-9 sig figs
• Multiplying dividing is a lot more common,
especially since conversions involve these
• To maintain the number of significant
figures you count the number of significant
figures for each thing being multiplied or
divided and the one with the least number is
how many your answer has
• Ex. 4.0 x 2.00 x 1.0000 = _________
1-9 sig figs
• Ex. 4.0 x 2.00 x 1.0000 = _________
• 4.0 has 2, 2.00 has 3 and 1.0000 has 5 sig
figs, so your answer should have 2
• 4.0 x 2.00 x 1.0000 = 8.0000 = 8.0
• Exceptions, exact numbers have infinite sig
figs
• Ex. Counting numbers (2 people),
conversion factors (3 feet = 1 yard)
1-9 sig figs
• Rounding sig figs to X places
• To round 1.2345678 to 4 places, you go to
the 4th place (the 4), then if the number
directly to the left is 5 or larger you round
up, if it is 4 or lower you round down
• 1.235
• Significant figures and significant digits are
the same thing and will be used
interchangeably as will sig figs and sig digs
1-9 sig figs
• The most important things about sig figs at
this point is understanding why we use
them, and being able to correctly count 0’s
as significant or not. The rest will get lots
of practice.
1-10 The unit factor method
• This chapter is devoted to conversions,
there are two ways to understand
conversions
– I think of conversions as multiplying by one,
you take some value, and multiply it and divide
by two equivalent values that are different units
– You can also understand a conversion as a
process and use matching units and routine to
do them
1-10
• Ex. Convert 24 inches into feet
• 24 inches * (1 foot) = 2 feet
(12 inches)
• Because 1 foot = 12 inches, you have
changed the units without changing the
value of it
1-10
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•
•
•
•
Ex. Convert 2 miles into feet
Step 1 write down given
2 miles
Step 2 make the crossbow
2 miles * (
)
(
)
1-10
• Step 3 fill in units of given on bottom and
desired units on top
• 2 miles * (
feet )
(
miles )
• Step 4 Fill in equivalent amounts
• 2 miles * ( 5280 feet ) = 10,560 feet
( 1 mile )
1-11 Percentage
• % by mass is used in chemistry and it is the
same as the rest of the percentage
calculations you’ve done previously, it’s
just more intimidating because it involves
chemicals
• Always remember that percents are a part of
a whole, so you will always have a part, and
the whole.
1-12 Density and specific gravity
• Density is mass/volume, it is an intensive
property so it does not depend on amount
• Density makes a good conversion factor
between volume and mass
• Specific density is the density of an object
relative to the density of water
• To get specific density you just remove the
units (or divide by the density of water)
1-13 Heat and temperature
• Converting between fahrenheit and celsius
is almost never needed and can be looked
up when it is, however, Celsius to Kelvin
conversions occur all of the time
• Since the lowest temperature possible is 0
K, and Celsius can be negative, to get from
Kelvins to Celsius you subtract 273 and add
273 to get from Celsius to Kelvin
1-13
• You should also know the freezing and
boiling points of water in all three scales
• Fahrenheit 32 and 212
• Celsius 0 and 100
• Kelvin 273 and 373
• Kelvins do not include the degree symbol
1-14 heat transfer and the
measurement of heat
• Joules are the most common units of energy
• Calories are also used, a calorie is 4.184 J,
and a Calorie or kilocalorie is 4184 J. The
Calorie with a big C is the kind you see on
food packaging
• Specific heat capacity is the amount of heat
it takes to heat up 1 gram of a substance
1-14 Specific heat capacity
• Q = mCt where Q is heat (J), m is mass
(g), C is a specific heat constant (J/(g °C),
and t is the change in temperature (°C)
• C depends on the material, for water C is
4.184 J/(g °C) and for metals the specific
heat is much lower because they heat up to
higher temperatures with less heat
1-14
• Heat capacity is like specific heat capacity
only it is for an entire object, not 1 gram of
a substance. Heat capacity is often used for
calorimeters and bomb calorimeters and
when the mass is unknown
• A calorimeter is a device that usually
consists mostly of water that the
temperature change can be related to find
the amount of heat absorbed or released
1-14
• Sample calorimetry problem