Honors Chemistry



Chapter 2 problem set:

Page 42 Section Review # 3, 5

Page 57 Section Review # 2, 5, 7

Page 59-61 Chapter Review # 15, 16, 18, 19, 22, 24, 25, 29, 35, 36, 38, 49, 51

2.1 Scientific Method (see text)

2.2 Units of Measurement

Measurement – _______________________________________________________________________

The problem is, what do you use as a standard?

Standard should be an ___________________________________________________________________________________________________________________________________________________________.

The SI System

Important base units to know:

|Quantity |Unit |Abbreviation |

|Mass |kilogram | |

|Length |meter | |

|Time |second | |

|Temperature |kelvin | |

|Amount of substance |mole | |

|Electric current |ampere | |

Important prefixes(multiples of base units) to know:

|Prefix |Abbreviation |Meaning |Example |

|mega- |M |106 | |

|kilo- |k |103 | |

|deci- |d |10-1 | |

|centi- |c |10-2 | |

|milli- |m |10-3 | |

|micro- |u |10-6 | |

|nano- |n |10-9 | |

|pico- |p |10-12 | |

Factor Label Method (Dimensional Analysis)

A method of problem solving that treats units like algebraic factors

Rules

1. Put the known quantity over the number 1.

2. On the bottom of the next term, put the unit on top of the previous term.

3. On top of the current term put a unit that you are trying to get to.

4. On the top and bottom of the current term, put in numbers in order to create equality.

5. If the unit on top is the unit of your final answer, multiply/divide and cancel units. If not, return to step # 2.

6. As far as sig figs are concerned, end with what you start with!

Example: convert 3.0 ft to inches

Example: convert 1.8 years to seconds

Example: convert 2.50 ft to cm if 1 inch = 2.54 cm

Example: convert 75 cm to m

Example: convert 150 g to kg

Example: convert 0.75 L to cm3

.

Example: convert 22 cm to dm

Density – ratio of mass to volume

The common density units are:

•

•

•

Formula is _________________________________

Density is a) b) c)

Two ways to find volume in density problems:

1. 2.

Note: the density is the same no matter what is the size or shape of the sample.

Ex1: Find the density of an object with m= 10g and v=2 cm3

Ex2: A cube of lead 3.00 cm on a side has a mass of 305.0 g. What is the density of lead?

First, calculate it’s volume:

Next, calculate the density:

2.3 Using Scientific Measurement

Precision vs. Accuracy

|Precision |Accuracy |

| | |

| | |

|_____________________________ |[pic] |______________________________ |

| | | |

| | | |

| | | |

|_____________________________ | |______________________________ |

Percent Error - experiments don’t always give true results - error is pretty much a given

Observed value (experimental value) - data found in an ____________________

True value (accepted value, theoretical value) - data that is generally accepted as true

Percent (%) error = (order is important as it implies direction)

+/- shows show the direction of the error - values are either too high or too low

Note: Percent error is a positive number when the experimental value is too high and is a negative number when the experimental value is too low.

Ex1: 66 Co is the answer in your experiment 65 CO is the theoretical value

Uncertainty in Measurement

• Two important points:

1. ______________________________________________________________________________

2. ______________________________________________________________________________

• Example – your chemistry’s teacher’s height is 5’4” not 5.43234566333333 inches

• When making a measurement, include all readable digits and 1 estimated digit

o always read between the lines!

o the digit read between the lines is always uncertain

|[pic] |if the measurement is exactly half way between lines record it as 0.5 |

| |if it is a little over, record ___________________ |

| |if it is a little under, record __________________ |

| |You would read this as 18.0 mL and not 18.5 mL. |

Significant Figures (Digits) - “Sig Figs”

• Definition: _____________________________________________________________________

• 1.15 ml implies 1.15 + 0.01 ml

• The more significant digits, the more reproducible the measurement is.

These are the numbers that “count!”

Ex1: π = 22/7 = 3.1415927 what do math teachers let you use?

Ex2: You collect a paycheck for a 40 hour week – what’s the difference between getting paid pi vs. 3.14 ?

Rules for finding the # of sig figs

1. All non-zeros are significant ex. 7 [ ] 77 [ ] 4568 [ ]

2. Zeros between non-zeros are significant ex. 707 [ ] 7053 [ ] 7.053 [ ]

3. Zeroes to the left of the first nonzero digit fix the position of the decimal point and are not significant

ex: 0.0056 [ ] 0.0789 [ ] 0.0000001 [ ]

4. In a number with digits to the right of a decimal point, zeroes to the right of the last nonzero digit are significant

ex: 43 [ ] 43.00 [ ] 43.0 [ ] 0.00200 [ ] 0.040050 [ ]

5. In a number that has no decimal point, and that ends in zeroes (ex. 3600), the zeroes at the end may or may not be significant (it is ambiguous). To avoid ambiguity, express in scientific notation and show in the coefficient the number of significant digits. ex. 3600 = 3.6 x 103 [ ]

Counting significant digits

• 1. __________________________________________________________

• 2. __________________________________________________________

Ex1: 700 [ ] - means “about 700 people at a football game”

700. [ ] - means “exactly 700 ......”

700.0 [ ] - means “teacher weighs exactly 700.0 lbs”

Other examples 0.5 [ ] 0.50 [ ] 0.050 [ ]

Sig. figs apply to scientific notation as well

9.7 x 10 2 = 970 [ ] 1.20 x 10 -4 = .000120 [ ]

Calculating with Measurements ( Sig Fig Math )

Rounding Rules

XY ---------------------> Y

_____________________________________________________________________

_____________________________________________________________________

_____________________________________________________________________

_____________________________________________

_____________________________________________

Ex1: round to 3 sig figs 35.27 =

87.24 =

95.25 =

95.15 =

Note - the “5” rule only applies to a “dead even” 5 - if any digit other than 0 follows a 5 to be rounded, then the number gets rounded up without regard to the previous digit.

Ex2: round to 3 sig figs 35.250000000000000000000000001 =

Calculating rules:

1. Multiplying or dividing – round results to the smaller # of sig. figs in the original problem.

Ex1: 3.10 cm Ex2: 7.9312 g

4.102 cm / 0.98 m

x 8.13124 cm

2. Adding or subtracting - round to the last common decimal place on the right.

Ex1: 21.52 Ex2: 73.01234 g

+ 3.1? - 73.014?? g

Note - exact conversion factors do not limit the # of sig figs - the final answer should always end with the # of sig figs that started the problem

ex. convert 7866 cm to m

2.4 Solving Quantitative Problems

1. Analyze – read carefully, list data with units, draw a picture

2. Plan – list conversion factors, show that units will work

3. Compute – use a calculator and use significant figures

4. Evaluate – does the answer “seem right” ?

Proportional relationships

Directly proportional – examples – density (mass of water vs. volume of water), grades vs. freedom

[pic]

In this example, we would say that, “volume of water is directly proportional to mass of water.” We can write it as __________________________

Inversely proportional – examples – speed vs. time, more accidents = less driving

Another chemistry example, as the pressure of a gas increases, the volume decresases:

|Volume of |[pic] |

|gas (mL) | |

|Pressure | |

|of gas (atm) | |

| | |

|25.2 | |

|0.971 | |

| | |

|28.3 | |

|0.865 | |

| | |

|32.4 | |

|0.755 | |

| | |

|39.6 | |

|0.618 | |

| | |

|43.1 | |

|0.568 | |

| | |

|47.8 | |

|0.512 | |

| | |

|52.6 | |

|0.465 | |

| | |

When two variables are related this way, they are said to be inversely proportional.

We can write it as __________________________

2.2 – Heat and Temperature – there is a difference

Compare heat and temperature of a thimbleful of water vs. a bathtub.

Heat is the amount of energy a chemical has –measure it in joules (J). Because we can’t directly measure heat, we have to measure “temperature”, which reflects how much kinetic energy an object has – measure it in °C or Kelvins). In thermodynamics, the term “enthalpy” is used interchangeably with “heat”, as it avoids confusion between the terms “heat” and “temperature”.

Heat vs. Temperature

Temperature – a measure of the average kinetic energy of the particles

Heat – the sum of the total energy of the particles

Temperature scales

Thermometer – works by thermal expansion, usually Hg or alcohol

Celsius - devised by a Swedish astronomer - Anders Celsius - 1742

Kelvin - named for Lord Kelvin - English physicist

Absolute zero - temperature at which all molecular motion stops - never been reached, temperature = 0 K or -273.15 oC

Conversion formulas:

Conversion formulas:

K =

oF =

oC =

Examples:

Ex1 Convert 100.0 oC to K

Ex2 Convert 293.0 K to oC

Ex3 Convert 90.0 oF to oC

Ex4 Convert 20.0 oC to oF

Heat transfers between objects – flows from hot to cold - Law of Conservation of Energy

Ex1:ice cube in a thermos of hot water - ice melts, water cools - same amount of heat

SI unit of heat - Joule (J) calorie is also used frequently

Calorie - the amount of energy required to raise the temperature of 1 g of water by 1 oC

(Calories – capital letter – really means kilocalories – used in food energy measurement)

1.000 calorie = 4.184 Joules

Heat Capacity vs. Specific Heat

Amount of heat needed to change the temperature of an object depends on:

a) b) c)

Heat capacity - amount of heat needed to raise the temperature of a sample – depends on mass – we need something more specific, so…..

Ex, which would you rather use to pull a pan from a hot oven, an oven mitt or a sheet of aluminum foil? The aluminum foil will transmit the heat easily while the oven mitt is a much better insulator. The reason: Oven mitts have a higher heat capacity than aluminum.

Specific Heat ( Cp ) - the amount of energy required to raise the temperature of 1 g of substance by 1 Co

q = heat (J or cal)

Cp = specific heat (J / g ▪ °C) or (cal / g ▪ °C)

m = mass(g)

Δt = change in temperature (°C or K)

so.... Cp = 1.000 cal/g oC or 4.184 J/g oC for water

Different materials have different specific heats

|material at 298 K and 1 atm |Cp specific heat (J/g K) |

|ice |2.09 |

|water |4.18 |

|steam |1.86 |

|sodium |1.23 |

|aluminum |0.9 |

|iron |0.45 |

Computing heat to determine how much heat is required to heat a material

It takes energy to make the temperature of anything increase. The relationship between energy and temperature is shown by the equation:

Specific Heat Problems

For water, Cp = 1.000 cal/g oC or 4.184 J/g oC for water

Ex1: How many calories does it take to heat 20. g of water from 10.0 to 40.0 oC? Also how many J?

Ex2: How much heat is required to heat 75 g of Iron (Cp = 0.444 J/gCo) from 15.5 to 57.0 oC?

Ex3: What is the specific heat of an object if 250 calories will heat 55 g of it from 25.0 to 100.0 oC?

Ex4: - If a 100.0 g sample of silver (Cp = .237 J/g oC) at 80.0 Co loses 50. calories, what will its final temperature be?

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