CHEM 1311/1111: Exam 1 - Ch. 1 - Phases and Classification of Matter

The Scientific Method

A method of acquiring knowledge and testing it.

Hypothesis

A tentative explanation of observations that act as a guide for gathering and checking information.

Law

A summary of a large number of experimental observations.

Theory

An explanation of observations that can be tested and it’s well substantiated.

Hypothesis - A testable, educated prediction about what might happen.

ex) If I increase the temperature of a gas, then its pressure will increase.

Law - A description of a pattern in nature that consistently occurs, often expressed mathematically.

ex) Newton’s Law of Gravitation describes how objects attract each other.

Theory - A well-supported explanation for why or how something happens, backed by lots of evidence.

ex) Cell Theory explains that all living things are made of cells.

AN EASY WAY TO REMEMBER:

Hypothesis = “What do I think will happen?”

Law = “What happens?”

Theory = “Why/how does it happen?”

What is the difference between a hypothesis, law, and theory?

NOTES

Solids

Densely packed, close contact between the molecules, little to no movement. Mostly vibrational motion. Molecules are stuck in place.

Liquids

Dense, but less dense than solids. More movement between the molecules. Molecules are free to move around past each other.

Gases

Are not dense at all, with mostly “empty” space. Atoms and molecules expand to fill the available space. Molecules and atoms in the gas phase bounce off of each other.

Matter

A substance that has mass and takes up space by having volume.

Mass

How much of a substance there is.

Volume

The amount of space that a substance has. As on the cube below, height (h), times width (w), times depth (d)

Matter

________ can’t be created or destroyed. This is called the Law of Conservation of Mass.

Atoms (or elements)

The basic chemical unit of matter. Indivisible without changing the nature of the element.

Molecules

Atoms connected through a “chemical bond” (such as covalent or ionic). These could be atoms of the same element or different elements. For example, H2, H2O, O2, H2O2.

-The chemical bond is made up of electrons.

Compounds

Similar to molecules but made up of different elements. For example, H2O, CO2, CO, NaCl.

Heterogeneous

Composition varies through the mixture.

Example) Salt and sugar mixed as solids.

Homogeneous

Composition is the same throughout the mixture.

Example) Salt in water or Sugar in water

NOTES

Characteristic properties of a substance can be observed and measured, such as color, phase at a given temperature, boiling point, melting point, density.

Density

The amount of mass in a given volume.

Extensive properties

Depend on the amount of substance.

Example) Mass (grams), volume (in liters), etc.

Intensive Properties

Do NOT depend on amount of substance.

Example) Temperature, density, boiling point, freezing point, color.

Basic Unit: Meter (m)

Property: Length

Basic Unit: Grams (g)

Property: Mass

Basic Unit: Seconds (s)

Property: Time

Basic Unit: Kelvin (K), also Fahrenheit (F) and Celsius (C)

Property: Temperature

Basic Unit: Liter (L)

Property: Volume

1 kilometer (Km)

1,000m = ________________

1 decimeter (dm)

0.1m = ___________________

1 centimeter (cm)

0.01m = __________________

1 millimeter (mm)

0.001m = _________________

1 micrometer (µm)

0.000001m = _____________________

NOTES

Another form of expressing volume is cubic centimeters (cm^3).

As on the example above of a cube in which the height, width, and depth are multiplied, if these are measured in cm, the answer will be in units of cm^3.

Therefore, 1 cm^3 = 1 mL

NOTES

Density (d) - The amount of mass in a given volume. (D=M/V)

For example, if a cube is 30 mL in volume and its mass is 25 grams, its density would be 0.833 g/mL.

25/30 = 0.833

All substances have their characteristic densities:

H2O: 1 g/mL

Iron: 7.9 g/mL

Copper: 9.0 g/mL

Since densities are ratio values, they can be used in calculations as conversion factors.

NOTES

°C = [°F – 32 °F][5 °C/9 °F]

°F = [9 °F/5 °C][°C] + 32 °F

K = °C + 273.15 (could also use 273)

Significant Figures

These figures tell the certainty of a measurement.

For example:

You place a substance on a balance to measure its mass, and it gives you a reading of 2.134419 g.

Not all numbers will be of value. To what certainty do you want the value to be represented?

The number could be rewritten as: 2.1, 2.13, 2.134, 2.1344, etc. depending on how many digits you want the accuracy of the number to be.

Rules for Determining Significant Figures in a number

• Any non-zero digit is significant.

• Zeros between non-zero digits are significant.

• Zeros to the left of non-zero digits are not significant.

• “Trailing” zeros at the end of a number are significant when the number is written showing a decimal point.

NOTES

Rounding up Numbers

If a measured number is 2.13765483, but you want to represent it with fewer significant figures, you may need to round up the number.

To properly do this, we look at the digit to the right of the last number we want to keep. If the digit to the right is less than 5, no changes are made. If the digit is 5 or larger, the last number we want to keep gets changed to the next higher number.

For example: 2.13765483

• Rounding up to two significant figures: 2.1

• Rounding up to three significant figures: 2.14

• Rounding up to four significant figures: 2.138

• Rounding up to five significant figures: 2.1377

NOTES

Scientific Notation

A specific way of writing numbers, especially for very large or very small numbers.

For small numbers:

Rewrite the number 0.0000000000213432 in scientific notation, rounding it up to four significant figures: 2.134 x 10-11

For large numbers:

Rewrite the number 5,014,387,917 in scientific notation rounding it up to four significant figures: 5.014 x 109

NOTES

IMPORTANT: THE ‘x’ DOES NOT REPRESENT A MULTIPLICATION!

The same rules for significant figures that we learned earlier apply to scientific notation, for the numbers that are to the left of the “x.”

Scientific notation can also be used for rewriting numbers with an ambiguous number of significant figures to rewrite them with accurate significant figures.

Example: 2,000 is ambiguous with its significant figures. We can rewrite it (as below) to represent sig. figs. better.

2.000 x 10^3 ← Now it has four significant figures.

2.00 x 10^3 ← Now it has three significant figures.

2.0 x 10^3 ← Now it has two significant figures.

NOTES

For addition and subtraction, the number of digits beyond the decimal point dictates the number of sig. figs.

ex) 12, 343.2 g + 0.1893 g = 12,343.3893 g

Answer with correct sig. figs.: 12,343.4 g

NOTES

For multiplication and division, we look at the whole number, and the number with the fewer digits dictates where to round up.

ex) 24,368.1 m x 55.384 m = 1,349,602.85 m2

Answer with correct sig. figs.: 1.3496 x 106 m2

Positive

Moving the decimal point to the left when converting a standard number to scientific notation makes the exponent __________.

Negative

Moving the decimal point to the right when converting a standard number to scientific notation makes the exponent __________.