Unit 1 Intro to Chemistry: Scientific Method, Metric System Classification of Matter, Changes and Significant Figures

Scientific Method

A systematic approach to problem-solving and answering questions through testing and observation.

Observation

Information gathered through the senses about a phenomenon or process.

Hypothesis

A testable prediction or tentative explanation for an observation.

Experiment

A procedure carried out to test a hypothesis under controlled conditions.

Scientific Theory

A well-substantiated explanation of some aspect of the natural world based on extensive evidence.

Scientific Law

A statement summarizing observed factual behavior in nature, often expressed mathematically.

Theory vs. Law

A theory explains why phenomena occur, whereas a law describes what happens.

Coke vs. Diet Coke Demo: Key Variable

Regular Coke contains sugar, making it denser than Diet Coke, which uses artificial sweetener (nutrasweet).

Metric System Base

A measurement system based on powers of 10.

Meter (m)

The metric base unit for measuring length.

Gram (g)

The metric base unit for measuring mass.

Liter (L)

The metric base unit for measuring liquid volume.

King Henry Mnemonic

King Henry Died Unexpectedly Drinking Chocolate Milk (Kilo, Hecto, Deka, Unit, Deci, Centi, Milli).

Kilo- (k)

Metric prefix representing 1,000 (10³).

Centi- (c)

Metric prefix representing 1/100 (10⁻²).

Milli- (m)

Metric prefix representing 1/1,000 (10⁻³).

Converting to a Smaller Unit

Multiply by 10 per step, moving the decimal to the right.

Converting to a Larger Unit

Divide by 10 per step, moving the decimal to the left.

Chemistry

The study of matter and the changes it undergoes.

Matter

Anything that has mass and volume.

Mass

The amount of matter in an object.

Energy

Anything that can do work or produce heat.

Weight

The force of gravity acting on an object's mass.

Solid (Properties)

Definite shape, definite volume, non-compressible, high density.

Liquid (Properties)

Indefinite shape, definite volume, non-compressible, less dense than solids.

Gas (Properties)

Indefinite shape, indefinite volume, compressible, very low density.

Pure Substance

Matter that has a uniform and definite composition.

Element

A pure substance found on the Periodic Table, represented by chemical symbols.

Compound

A pure substance formed when elements chemically combine in fixed ratios.

Mixture

Two or more pure substances physically mixed together without fixed ratios.

Heterogeneous Mixture

A mixture lacking uniform composition, where distinct parts can be visually identified.

Homogeneous Mixture

A mixture with uniform composition throughout, where individual parts cannot be visually distinguished.

Physical Property

A characteristic observed or measured without changing the composition or identity of a substance.

Chemical Property

A property indicating how a substance reacts with other substances to form new matter.

Physical Change

A change in appearance or state of matter without altering its chemical composition.

Chemical Change

A process where starting substances react to form new substances with distinct properties.

Examples of Physical Changes

Cutting, breaking, pulverizing, melting, freezing, boiling, and subliming.

Examples of Chemical Changes

Rusting, burning, corrosion, digestion, respiration, and decaying.

Reactants

The starting substances in a chemical reaction, written on the left side of the arrow.

Products

The new substances formed by a chemical reaction, written on the right side of the arrow.

Law of Conservation of Mass

Matter cannot be created or destroyed in any physical or chemical change.

Measurement

A type of observation.

Qualitative Measurement

A descriptive observation without numbers, such as color, temperature feel, or texture.

Quantitative Measurement

An observation made with an instrument that includes both a number and a unit.

Accuracy

How close a measurement is to the true or accepted value.

Precision

How close multiple measurements are to each other.

Accuracy vs. Precision Mnemonic

ACcurate = Correct; PRecision = Reproducibility.

Scientific Notation

A shorthand method for writing very large or very small numbers.

Parts of Scientific Notation

A coefficient between 1 and 10 multiplied by a power of 10.

Scientific Notation Exponent Rule

Numbers greater than 1 use positive exponents; numbers less than 1 use negative exponents.

Calculator Key for Scientific Notation

Use the "EE" or "EXP" button instead of typing "× 10^".

Precise but Not Accurate Measurement Example

Getting 5.0g, 5.1g, and 5.0g on a balance when the true mass is 10.0g.

10,000 in Scientific Notation

1 × 10⁴

0.000157 in Scientific Notation

1.57 × 10⁻⁴

3.5 × 10⁻⁵ in Long Form

0.000035

7.1 × 10⁴ in Long Form

71,000

Product of (6 × 10⁻⁸) × (4 × 10⁻⁴)

2.4 × 10⁻¹¹

Product of (7 × 10⁶) × (8 × 10⁵)

5.6 × 10¹²

0.00304 in Scientific Notation

3.04 × 10⁻³

503 in Scientific Notation

5.03 × 10²

2,000,000 in Scientific Notation

2 × 10⁶

(6.02 × 10²³) / (1.24 × 10²⁴)

0.485

(6 × 10⁹) × (1.35 × 10¹⁸)

8.1 × 10²⁷

Measurement Reading Rule

Includes all certain digits read directly off the device plus one estimated digit.

Exact Numbers (Sig Figs)

Numbers with no uncertainty (e.g., counting objects or defined quantities) have infinite significant figures.

Sig Fig Rule: Non-Zero Digits

All non-zero digits are always significant (e.g., 2.17 has 3 sig figs).

Sig Fig Rule: Leading Zeros

Zeros to the left of non-zero digits are never significant (e.g., 0.003 has 1 sig fig).

Sig Fig Rule: Captive Zeros

Zeros between non-zero digits are always significant (e.g., 205 has 3 sig figs).

Sig Fig Rule: Trailing Zeros (With Decimal)

Zeros at the end of a number ARE significant if a decimal point is present (e.g., 155.0 has 4 sig figs).

Sig Fig Rule: Trailing Zeros (Without Decimal)

Zeros at the end of a number are NOT significant if there is no decimal point (e.g., 1550 has 3 sig figs).

Significant Figures in 23

2

Significant Figures in 1.498

4

Significant Figures in 248.0

4

Significant Figures in 76000

2

Significant Figures in 32.80

4

Significant Figures in 0.238

3

Significant Figures in 80.035

5

Significant Figures in 0.055870

5

Significant Figures in 0.00450

3 significant figures

Significant Figures in 1200

2 significant figures (unless a decimal point is present)

Significant Figures in 6.02 × 10²³

3 significant figures

Significant Figures in 0.070

2 significant figures

Significant Figures in 100.0

4 significant figures

Sig Fig Addition/Subtraction Rule

Round the answer to the least number of decimal places.

Sig Fig Multiplication/Division Rule

Round the answer to the least number of significant figures.

12.11 + 0.3 (Sig Figs)

12.4 (rounded to 1 decimal place)

4.56 × 1.4 (Sig Figs)

6.4 (rounded to 2 sig figs)

8.0 ÷ 2.00 (Sig Figs)

4.0 (rounded to 2 sig figs)

15.678 - 2.1 (Sig Figs)

13.6 (rounded to 1 decimal place)

0.00560 × 300 (Sig Figs)

1.7 (rounded to 2 sig figs)

Why Scientists Use Significant Figures

To reflect measurement precision and avoid overrepresenting accuracy beyond tool limits.

5.7 m + 2.11 m (Sig Figs)

7.8 m (rounded to 1 decimal place)

7.12 g - 1.345 g (Sig Figs)

5.78 g (rounded to 2 decimal places)

100.45 mL - 2.0 mL (Sig Figs)

98.5 mL (rounded to 1 decimal place)

17.85 g / 2.13 mL (Sig Figs)

8.38 g/mL (rounded to 3 sig figs)

0.0561 cm × 0.030 cm (Sig Figs)

0.0017 cm² (rounded to 2 sig figs)

4 cm × 2.3 cm (Sig Figs)

9 cm² (rounded to 1 sig fig)

11.7 cm × 3.522 cm × 0.013 cm (Sig Figs)

0.54 cm³ (rounded to 2 sig figs)

Round 345 to 2 significant figures

350

Round 7850 to 1 significant figure

8000