Chapter 1: Chemistry Basics – Matter and Measurement

Chemistry Basics: Matter and Measurement

Classifying Matter: Pure Substance or Mixture

Matter is anything that has mass and occupies space. It can be classified based on its composition as either a pure substance or a mixture.

• Pure Substance: A form of matter with a constant composition and distinct properties. Examples include elements (e.g., oxygen, gold) and compounds (e.g., water, citric acid).

• Mixture: A combination of two or more substances in which each retains its own identity. Mixtures can be homogeneous (uniform composition, e.g., saltwater) or heterogeneous (non-uniform, e.g., salad).

Elements and the Periodic Table

Elements are pure substances that cannot be broken down into simpler substances by chemical means. They are organized in the periodic table based on their atomic number and properties.

• Element Symbol: Each element is represented by a unique one- or two-letter symbol (e.g., H for hydrogen, O for oxygen).

• Groups and Periods: Vertical columns are called groups; horizontal rows are periods.

Elements and Nutrition

Certain elements are essential for human nutrition and biological processes. For example, sodium (Na), potassium (K), calcium (Ca), and iron (Fe) are vital for health.

Elements in Compounds

Atoms of different elements can combine to form compounds. The chemical formula shows the types and numbers of atoms present.

• Example: Citric acid (C6H8O7) contains 6 carbon, 8 hydrogen, and 7 oxygen atoms.

Physical Changes

A physical change alters the form or appearance of matter but does not change its chemical identity.

• Examples: Melting ice, boiling water, dissolving sugar in water.

Chemical Changes

A chemical change results in the formation of one or more new substances with different chemical identities. This process is called a chemical reaction.

• Example: Burning wood, rusting iron.

Chemical Equations

Chemical equations represent chemical reactions, showing reactants and products, their physical states, and the direction of the reaction.

• General Format:

$\mathrm{Reactant\ (state)\ +\ Reactant\ (state)\ \rightarrow\ Product\ (state)}$

• State Symbols: (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous (dissolved in water)

How to Balance Chemical Equations

Balancing equations ensures the same number of each type of atom on both sides, reflecting the Law of Conservation of Mass.

• Example:

$\mathrm{HCl(aq)\ +\ Zn(s)\ \rightarrow\ ZnCl_2(aq)\ +\ H_2(g)}$

Measurement Systems and SI Units

Measurements in science use defined systems for consistency. The Système International d’Unités (SI) is the modern metric system.

• Mass: kilogram (kg)

• Volume: liter (L)

• Length: meter (m)

Prefixes modify units by powers of ten (e.g., milli-, centi-, kilo-).

Metric Prefixes

Prefix	Symbol	Factor
kilo-	k	103
centi-	c	10-2
milli-	m	10-3
micro-	μ	10-6
nano-	n	10-9

Conversion Factors

Conversion factors are ratios used to express a quantity in different units.

• Example: $1\ \mathrm{kg} = 1000\ \mathrm{g}$

Significant Figures (Sig Figs)

Significant figures reflect the precision of a measured quantity. Only measured values (not exact numbers) are considered for sig figs.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if there is a decimal point.

Calculating with Significant Figures

• Addition/Subtraction: The result should have the same number of decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

Scientific Notation

Scientific notation expresses very large or small numbers as a product of a number between 1 and 10 and a power of ten.

• Example: $3.25 \times 10^2\ \mathrm{mg}$

Practice Problem Example

• Problem: Convert 325 mg, 303.15 mg, and 418.1662 mg to grams, add, and report with correct sig figs in scientific notation.

• Solution:

  • 325 mg = 0.325 g

  • 303.15 mg = 0.30315 g

  • 418.1662 mg = 0.4181662 g

  • Total = 1.0463162 g (report as 1.05 g or $1.05 \times 10^0$ g, depending on sig figs)

Density

Density is the ratio of mass to volume. It is a physical property used to identify substances and as a conversion factor.

$d = \frac{m}{V}$

• Units: g/mL or g/cm3

Specific Gravity

Specific gravity is the ratio of the density of a substance to the density of water (at 4°C). It is unitless.

$\text{Specific gravity} = \frac{\text{density of substance}}{\text{density of water}}$

Energy

Energy is the capacity to do work or supply heat. It exists as potential (stored) or kinetic (motion) energy.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

• SI Unit: joule (J)

• Calorie: 1 cal = 4.184 J

• Nutritional Calorie: 1 Cal = 1000 cal

Heat and Specific Heat

Heat is energy transferred due to temperature difference. Specific heat is the amount of heat required to raise the temperature of 1 g of a substance by 1°C.

$q = m \cdot c \cdot \Delta T$

• Where q = heat (J), m = mass (g), c = specific heat (J/g·°C), ΔT = temperature change (°C)

States of Matter

Matter exists in three main physical states: solid, liquid, and gas.

• Solid: Definite shape and volume

• Liquid: Definite volume, takes shape of container

• Gas: No definite shape or volume

Properties of Different States of Matter

State	Shape	Volume	Compressibility
Solid	Definite	Definite	Very low
Liquid	Indefinite	Definite	Low
Gas	Indefinite	Indefinite	High

Accuracy and Precision

Accuracy refers to how close a measurement is to the true value. Precision refers to how close repeated measurements are to each other.

• Example: If a scale gives the same reading every time (precise), but it is not the true weight (not accurate).

Units and Dosing in Health Care

Health professionals use SI/metric units and must also understand U.S. customary units for dosing and reporting.

Reading Lab Reports

Lab reports use standard units, significant figures, and scientific notation to communicate results clearly and accurately.

Calculating a Dosage

• Example: Drug X: 5 mg/kg/day in two divided doses. Patient weighs 44 lb.

• Steps:

  1. Convert weight to kg: $44\ \mathrm{lb} \times \frac{1\ \mathrm{kg}}{2.205\ \mathrm{lb}} = 19.95\ \mathrm{kg}$

  2. Total daily dose: $19.95\ \mathrm{kg} \times 5\ \mathrm{mg/kg} = 99.75\ \mathrm{mg}$

  3. Per dose: $99.75\ \mathrm{mg} \div 2 = 49.88\ \mathrm{mg}$

Calculating Drop Units (IV Drip Rate)

• Example: 2.0 L fluid over 18 h, drop factor 20 gtt/mL. Calculate gtt/min.

• Steps:

  1. Convert volume to mL: $2.0\ \mathrm{L} = 2000\ \mathrm{mL}$

  2. Total time in minutes: $18\ \mathrm{h} \times 60 = 1080\ \mathrm{min}$

  3. Flow rate: $\frac{2000\ \mathrm{mL}}{1080\ \mathrm{min}} = 1.85\ \mathrm{mL/min}$

  4. Drip rate: $1.85\ \mathrm{mL/min} \times 20\ \mathrm{gtt/mL} = 37\ \mathrm{gtt/min}$

Percents in Health

• Example: Adult dose for sodium thiopental is 280 mg. A 19-lb child should receive 17% of the adult dose.

• Calculation: $280\ \mathrm{mg} \times 0.17 = 47.6\ \mathrm{mg}$

Additional info: Some calculations and tables were expanded for clarity and completeness.