HN Chemistry Unit 1
Chapter 1
chemistry – the study of matter and how matter changes
matter – anything that has mass and volume
pure chemistry – pursuing scientific, specifically chemical, knowledge for its own sake
For example:
What is in an atom?
How old is the universe?
applied chemistry – using scientific knowledge, specifically chemical knowledge for a practical application
For example:
Developing a cure for Ebola
Designing an efficient battery
hypothesis – a possible explanation for an observation that needs to be tested
An experiment – a method for testing a hypothesis, by changing only one variable at a time
manipulated (or independent) variable – the variable that is deliberately changed in an experiment
- When graphed, this variable is graphed on the x-axis.
responding (or dependent) variable – the variable observed during an experiment
- When graphed, this variable is graphed on the y-axis.
When graphing, both axes must be labelled and include units. The graph should be titled according to the variable on the y-axis versus the variable on the x-axis (y vs. x). A graph with a horizontal slope of about zero shows no relationship between the variables. When a graph indicates a relationship between the variables, linear data points can be used to determine the slope of the graph.
theory – a well-tested explanation for a wide-ranging series of observations
scientific law – a statement summarizing the results of extensive measured experiments and observations
Chapter 2
Properties of matter
mass – the measure of the amount of matter that an object contains
volume – the amount of space occupied by a sample of matter
extensive properties – depend on how much matter is present (for example: mass, volume)
intensive properties – depend on the identity of the matter present (for example: hardness, boiling point, etc.)
physical property – a condition or quality of a substance that can be observed without changing the composition of the substance
solids: particles are in close contact with each other in fixed positions, not compressible, shape and volume are defined
liquids: particles are in contact with one another but are free to move past one another, positions are not fixed, not particularly compressible, volume is defined but shape is not
gases: particles are very far apart, compressible, neither volume nor shape are defined
vapor: the gaseous state of a substance that is usually not a gas at room temperature
physical change – the composition of matter remains constant
Reversible – melting/freezing, evaporating/condensing
Irreversible – cutting, cracking
Mixtures
heterogeneous – a non-uniform mixture (consisting of at least two phases) in which the components are not evenly distributed
Example: chocolate chip cookies ☺
homogeneous (also known as a solution) – a uniform mixture (consisting of only one phase) in which the components are evenly distributed
Example: Gatorade
phase – part of a sample with a uniform composition and properties
Using physical properties to separate mixtures:
filtration – solid particles are separated from a heterogeneous mixture
distillation – using the boiling point of a substance can separate it from a different component of the mixture, the substance can be vaporized and then condensed in a separate container
Elements and Compounds
Matter can be classified as either:
- a pure substance with a definite chemical composition (chemical formula)
elements can’t be chemically decomposed into simpler substances For example: Na, K, Cl, O
o compounds can be chemically decomposed into simpler substances (elements) For example: H2O, CO2, CH4
or
- a mixture of pure substances with a composition that varies
o heterogeneous (see above)
o homogeneous (see above)
chemical change (chemical reaction) – the chemical configuration of matter changes
Example: 6CO2 + 6H2O → C6H12O6 + 6O2 (Carbon dioxide and water transform into glucose and oxygen.)
chemical symbol – the one or two printed letters used to represent an element
(The first letter is capitalized; the second letter must not be.) For example: Na, not NA, not NA, not Na.
The Periodic Table
- elements are arranged in increasing order of atomic number
- horizontal rows are called periods
- vertical columns are called groups (or families)
Chemical Reactions
chemical property – a condition or quality of a substance that can only be observed by changing the composition of the substance
reactants – chemicals present before a reaction (such as CO2 and H2O above)
products – chemicals produced by the reaction (such as C6H12O6 and O2 above)
Evidence that suggests a chemical reaction is taking place:
- a gas may form
- the color of a reactant may change
- energy may be given off or absorbed
- a precipitate may form (When two solutions are mixed, sometimes a solid forms as a result of a chemical reaction between the components of the solution. In this case, the solid is called a precipitate.)
Conservation of Mass – the total mass of products at the end of a reaction is the same as the total mass of all the reactants that
were present before the reaction
For example: If 264 grams of carbon dioxide react with 108 grams of water to produce 180 grams of glucose, then the mass of oxygen produced in the reaction must be 192 grams. The mass of the carbon dioxide and water equal 372 total
grams of reactants, so the total mass of the products must also equal 372 grams
Chapter 3
Measurements and Their Certainty
scientific notation – a method of expressing very large and very small numbers (know how to use scientific notation)
960,000 for example would be 9.6 x 10^5
0.000082 would be 8.2 x 10^–5
- The coefficient in scientific notation must be greater than 1, but it must be less than 10.
- A negative exponent indicates a relatively small number.
- A positive exponent indicates a relatively large number.
accuracy – how close is a measured value to the accepted value
precision – how close are the measurements to one another
error – the difference between the accepted value and the actual measurement
error = experimental value – accepted value
percent error – the percent by which a measured value differs from the accepted value
significant figures – reflect the certainty of a measurement; including all digits that are known and, in some cases, an additional estimated digit (know how to determine significant figures)
Rules for Significant figures
- all non-zero digits (1, 2, 3, 4, 5, 6, 7, 8, 9) ARE significant
- zeros (0) between non-zero digits ARE significant
- zeros that comes before (to the left of) any non-zero digits are NOT significant (They are placeholders and can are
eliminated when the number is expressed in scientific notation.)
- zeros after (→to the right of) non-zero digits are NOT significant when no decimal point is present (They are
placeholders and can be eliminated when the number is expressed in scientific notation.)
- zeros after (→to the right of) non-zero digits AND after (→to the right of) a decimal point ARE always significant
- a calculated answer CANNOT be more certain than the least certain measurement
- rounding in calculations depends on the type of calculation
o addition and subtracting (The final answer rounds to the fewest number of decimal places.)
o multiplication and division (The final answer rounds to the fewest number of significant figures.)
For some calculations, the final answer must be expressed in scientific notation to properly indicate the correct number of
significant figures.
For example:
12,500 mg/25 mL = 500 mg/mL, but this answer must be rounded to only 2 significant figures. 500 has 1
significant figure. 500. has 3 significant figures. The only way too round 500 to 2 significant figures is to express it in
scientific notation: 5.0 x 102 mg/mL
The International System of Units
Units of length: m, cm, km . . .
Units of volume: mL, L, m3, cm3 . . .
Units of mass: kg, g . . .
Mass and weight are not the same thing. Weight involves how much gravity is pulling on a given mass. Weight changes
with gravity, mass does not. A person weighs less on the moon, but has no less mass.
Units of temperature: K, oC, oF
absolute zero – zero on the Kelvin scale or –273.15oC
Units of energy: cal, J, kJ, kcal
density – the mass to volume ratio of an object (know how to use the density equation to solve either for density, mass, or (volume) density = mass
volume
Density can be affected by temperature. Matter usually expands when it warms. This can cause a substance to take up
more space at a greater temperature. Therefore, the calculated density would change. Lists of densities usually assume room
temperature.
Conversion Problems
conversion factor – equivalent measurements that can be used to convert between units (know how to convert between
units)
For example: 2.54 cm = 1 in or 10 mm = 1 cm or
dimensional analysis – method of cancelling units to solve problems (know how to solve a problem by cancelling units)
For example:
31.85 in x 2.54 cm x 10 mm x 1 atm = 1.064 atm