Chapter 1: Chemistry Basics – Matter & Measurement Study Notes

Chapter 1: Chemistry Basics – Matter & Measurement

1.1 – Classifying Matter: Pure Substance or Mixture

Chemistry is the study of matter and its changes. Matter is anything that occupies space and has mass. All matter can be classified as either pure substances or mixtures.

• Pure Substance: A sample containing only one type of substance, represented by a single chemical formula. Pure substances are either elements or compounds.

• Element: The simplest type of matter, composed of only one type of atom. Example: H (hydrogen)

• Atom: The smallest unit of matter that retains its chemical properties.

• Compound: A pure substance made of two or more elements chemically joined together. Example: H2O (water)

• Mixture: A combination of two or more substances that can be separated into their components.

• Homogeneous Mixture: Composition is uniform throughout. Example: Air

• Heterogeneous Mixture: Composition varies throughout. Example: Blood

Practice Problem 1.2

• Classify each as a pure substance or mixture: Well water, Ocean water, Cement, Tropical smoothie.

Practice Problem 1.4

• Classify mixtures as homogeneous or heterogeneous: Sports drink, Blueberry pancake, Gasoline, Raisin bran.

1.2 – Elements, Compounds, & the Periodic Table

The periodic table of elements organizes elements by their properties into periods (rows) and groups (columns).

• Groups: Vertical columns; elements in a group have similar chemical behaviors. Main group elements (1A–8A), transition elements (B groups).

• Periods: Horizontal rows, numbered 1–7. Periods 6 and 7 have sections set apart at the bottom.

• Boron Border: The staircase beginning at boron (B) separates metals from nonmetals. Elements bordering this line (except Al) are metalloids, which have properties of both metals and nonmetals.

• Chemical Symbol: One- or two-letter abbreviation for an element. Example: Hydrogen = H. Some symbols are derived from older languages.

• Chemical Formula: Shows the identity and number of elements in a compound. Subscripts indicate the number of atoms of each element. Example: H2O

Practice Problem 1.6

• Classify as element or compound: Fe, CaCl2, Si, KI.

Practice Problem 1.8

Name	Elemental Symbol	Group	Period	Metal or Nonmetal
Sodium	Na	1A	3	Metal
Phosphorus	P	5A	3	Nonmetal
Additional info: Table entries inferred for completeness.

Practice Problem 1.10

• Identify and count atoms in compounds: MgSO4, NaHCO3, C14H18N2O5.

1.3 – How Matter Changes

Matter can undergo physical or chemical changes. The state of matter refers to how a chemical physically exists: solid, liquid, or gas.

• Physical Change: Alters the state of matter without changing its chemical identity. Example: Melting ice (solid to liquid water)

• Chemical Change (Chemical Reaction): Involves the reaction of substances to form new substances; chemical identity is changed. Example: Burning charcoal

Practice Problem 1.12

• Determine if the following are physical or chemical changes: Making ice cream, Milk spoiling, Nail polish remover evaporating.

1.4 – Math Counts

Mathematics is essential for understanding chemistry, especially in measurement and calculations. The International System of Units (SI) is the preferred system for scientific measurement.

• SI Unit vs. Metric Unit: SI unit for mass is kilogram (kg); metric unit is gram (g).

• Precision: Relates to the instrument's ability to measure accurately, often reported to a specific number of decimal places.

• Significant Figures: Digits in a measurement known with certainty plus one estimated digit. Rules:

  • Significant if: Not zero, zero between nonzero digits, zero at end of number with decimal point.

  • Not significant if: Zero at beginning of decimal number, zero in large number without decimal point.

• Calculations:

  • Addition/Subtraction: Match least number of decimal places.

  • Multiplication/Division: Match least number of significant figures.

• Rounding:

  • If digit to be dropped is 4 or less, remove it.

  • If digit is 5 or greater, increase last retained digit by 1.

  • Do not round at each step in multi-step calculations; round only at the end.

• Percent (%): Part out of 100. Formula: $\text{Percent} = \text{fraction} \times 100$

• Convert fraction to percent: Divide numerator by denominator, multiply by 100, add % sign.

• Convert decimal to percent: Multiply by 100, add % sign.

Practice Problem 1.18

• Count significant figures in: 0.000068 g, 100 °F, 9,237,200 years, 25.00 m.

Practice Problem 1.20

• Perform calculations and report answers with correct significant figures.

Practice Problem 1.30

• Express numbers as percent: 0.58, 0.36, 0.125.

Practice Problem 1.32

• Express percent as fraction: 20%, 75%, 40%, 12%.

Practice Problem 1.34

• Determine percent from given numbers: Fifty is what percent of 125? Six is what percent of 600? etc.

1.5 – Matter: The “Stuff” of Chemistry

Matter is anything that occupies space and has mass. Key properties include mass, volume, and density.

• Mass: Amount of material in an object. Weight: Force on an object due to gravity.

• Volume: Space occupied by a substance. Clinical unit: Cubic centimeter (cc or cm3), 1 mL = 1 cc.

• Density (d): Ratio of mass (m) to volume (V). Formula: $d = \frac{m}{V}$

• Objects less dense than a liquid will float; more dense will sink.

• Specific Gravity (sp gr): Ratio of density of a substance to density of water. Formula: $\text{Specific gravity} = \frac{\text{Density of substance}}{\text{Density of water}}$

• Density of water is 1.00 g/mL at 4°C.

• Specific gravity in healthcare: Used to assess kidney function via urine density (normal range: 1.005–1.030).

• Temperature: Measure of hotness or coldness. Scales: Celsius (°C), Fahrenheit (°F), Kelvin (K). Conversions:

  • $°C = (°F - 32) \times \frac{5}{9}$

  • $°F = (°C \times \frac{9}{5}) + 32$

  • $K = °C + 273$

• Hyperthermia: Body temperature above 104°F. Hypothermia: Body temperature below 95°F.

Practice Problem 1.36

• Compare density of paper clip, table sugar, ice to liquid water.

Practice Problem 1.38

• Calculate density in g/mL for 2.0 L gasoline weighing 1.32 kg.

Practice Problem 1.39

• Interpret urine specific gravity of 1.002: dehydrated or overhydrated?

Practice Problem 1.42

• Convert -112°F to Kelvin.

States of Matter

The physical form in which a substance exists: solid, liquid, or gas.

• Solid: Particles in orderly arrangement, little motion; definite shape and volume.

• Liquid: Particles loosely associated, freely moving; definite volume, no definite shape.

• Gas: Particles not associated, rapidly moving; no definite shape or volume.

Property	Solid	Liquid	Gas
Shape	Definite	Adopts shape of container	Adopts shape of container
Volume	Definite	Definite	Indefinite
Particle Energy	Lowest	Intermediate	Highest
Particle Arrangement	Orderly, packed	Loosely packed, random	Far apart, random
Attraction Between Particles	Strong	Strong	Practically none

Practice Problem 1.50

• Identify state of matter: definite volume but takes shape of container; particles far apart, do not interact.

1.6 – Measuring Matter

Accurate and precise measurement is essential in chemistry and healthcare. Different systems of units are used, including SI, metric, and U.S. customary units.

• Accuracy: Measurements close to the true value.

• Precision: Repeat measurements are similar in value (reproducible).

• U.S. Customary System: Used in general American population; units include pound (lb), quart (qt), inch (in), etc.

Property	U.S. Customary Unit	SI/Metric Equivalent	U.S. Customary Equivalent
Mass	Pound (lb)	2.205 lb = 1 kg	16 oz = 1 lb
Volume	Quart (qt)	1.057 qt = 1 L	4 qt = 1 gal
Fluid Ounce (fl oz)	1 fl oz = 29.6 mL	8 fl oz = 1 cup
Teaspoon (tsp)	1 tsp = 4.93 mL	3 tsp = 1 tbsp
Length	Inch (in)	1 in = 2.54 cm	12 in = 1 ft

Practice Problem 1.54

• Convert 10 cc of medicine to teaspoons for dosing.

Practice Problem 1.56

• Calculate number of tablets for a child based on body weight and dosage.

IV Medicine Measurement

• Drop Factor (gtt/mL): Number of drops per milliliter; used to calculate drip rate for IV medicine.

• Drip Rate Formula: $\text{Drip rate (gtt/min)} = \frac{\text{Total volume (mL)} \times \text{Drop factor (gtt/mL)}}{\text{Time (min)}}$

Practice Problem 1.59

• Calculate drop rate for IV fluid delivery.

Practice Problem 1.60

• Calculate time to administer IV fluid at given drip rate.

Percent in Medicine and Nutrition

• Percent active ingredients: Used to determine dosage of potent medicines.

• Percent of adult dose: Children and pets may receive a percentage of the adult dose.

• Percent in nutrition labeling: Food items list percent daily values (%DV) for nutrients.

Practice Problem 1.61

• Calculate percent of active ingredient in a tablet.

Additional info: These notes cover foundational concepts in general chemistry, including classification of matter, the periodic table, physical and chemical changes, measurement, significant figures, density, specific gravity, temperature, and units of measurement. Practice problems are included to reinforce understanding and application of these concepts.