Introduction to Chemistry: Matter and Measurement (CH 117, Sacred Heart University)

Chapter 1: Chemistry Basics – Matter and Measurement

Learning Objectives

• Classify matter as pure substances or mixtures

• Distinguish between elements, compounds, and mixtures

• Understand the organization of the periodic table

• Describe physical and chemical changes

• Apply mathematical concepts in chemistry (significant figures, unit conversions, etc.)

• Measure and interpret mass, volume, density, and temperature

Classifying Matter

Pure Substances and Mixtures

Matter can be classified based on its composition:

• Pure Substance: Matter made up of only one type of substance, represented by a chemical formula or symbol.

• Element: The simplest type of matter, consisting of only one type of atom. An atom is the smallest unit of matter that retains its properties.

• Compound: A pure substance made of two or more elements chemically joined together.

• Mixture: A combination of two or more substances that can be separated by physical means.

• Homogeneous Mixture: Uniform composition throughout (e.g., salt water).

• Heterogeneous Mixture: Non-uniform composition; components are visibly distinct (e.g., salad, rocky road ice cream).

Example: Helium gas is a pure substance (element); cake batter is a mixture.

Elements, Compounds, and the Periodic Table

Structure of the Periodic Table

• Period: Horizontal row; corresponds to the number of energy levels (shells) in atoms.

• Group (Family): Vertical column; elements in a group have similar chemical properties and the same number of valence electrons.

Important Elements for Human Health: Elements such as carbon, hydrogen, oxygen, nitrogen, calcium, and phosphorus are most common in living things.

Compounds and Chemical Formulas

• A compound contains two or more elements chemically combined in a fixed ratio.

• A chemical formula shows which elements are present and how many atoms of each are in a compound.

Examples:

• Water: $H_2O$ (2 hydrogen, 1 oxygen)

• Table salt: $NaCl$ (1 sodium, 1 chlorine)

• Acetic acid: $C_2H_4O_2$ (2 carbon, 4 hydrogen, 2 oxygen)

• Phosphoric acid: $H_3PO_4$ (3 hydrogen, 1 phosphorus, 4 oxygen)

How Matter Changes

Physical and Chemical Changes

• Physical Change: Alters the form or appearance of matter but does not change its identity (e.g., melting, boiling, dissolving).

• Chemical Change: Alters the chemical identity of a substance; a new substance is formed (e.g., rusting, burning, digestion).

Examples:

• A copper penny turning green: chemical change

• Sugar melting: physical change

• Antacid tablet in water forming bubbles: chemical change

Chemical Equations and Balancing

• A chemical equation represents a chemical reaction, showing reactants and products.

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction; the mass of reactants equals the mass of products.

Example Equation:

• $C(s) + O_2(g) \rightarrow CO_2(g)$

Steps to Balance an Equation:

1. Examine the equation for balance.

2. Balance one element at a time using coefficients.

3. Check that all elements are balanced with the smallest whole-number coefficients.

Math Counts: Measurement and Calculations in Chemistry

SI Units and Prefixes

• The Système International d’Unités (SI) is the modern metric system.

• Prefixes (kilo-, centi-, milli-, etc.) change the unit by powers of 10.

• Equivalent units can be used as conversion factors (e.g., $1\,dL = 0.1\,L$).

Dimensional Analysis (Unit Conversions)

1. Determine the desired unit for the answer.

2. Identify the given information.

3. Choose conversion factors so units cancel, leaving only the desired unit.

4. Solve the problem.

General Formula:

• $\text{Given unit} \times \frac{\text{desired unit}}{\text{given unit}}$

Example: How many hours are in 0.5 years?

Significant Figures

• Significant figures (sig figs) reflect the precision of a measurement.

• Rules:

  • All nonzero digits are significant.

  • Leading zeros are not significant.

  • Zeros between nonzero digits are significant.

  • Trailing zeros are significant only if there is a decimal point.

Examples:

Measurement	Significant Figures
0.03 L	1
6.071 kg	4
20. g	2
12,000 km	2
3,450,000 m	3

Significant Figures in Calculations

• Addition/Subtraction: Answer should have the same number of decimal places as the measurement with the least decimal places.

• Multiplication/Division: Answer should have the same number of significant digits as the measurement with the least significant digits.

Rounding: If the digit to be dropped is 5 or greater, increase the last retained digit by 1.

Scientific Notation

• Expresses numbers as $C \times 10^n$, where $1 \leq C < 10$ and $n$ is an integer.

• Positive exponent: number greater than 1; negative exponent: number between 0 and 1.

• Only significant figures are shown in the coefficient.

Percentages

• Percent (%) means per 100.

• Formula: $\text{Percent} = \frac{\text{part}}{\text{whole}} \times 100$

Matter: The "Stuff" of Chemistry

Mass and Weight

• Mass: Measure of the amount of material in an object (unit: gram, g).

• Weight: Force of gravity on an object; can vary with location, but mass remains constant.

Volume

• Volume is the amount of space occupied by matter (unit: milliliter, mL; cubic centimeter, cm3).

• 1 mL = 1 cm3

Density and Specific Gravity

• Density (d): Ratio of mass to volume.

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

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

• Specific Gravity: Ratio of the density of a sample to the density of water (unitless).

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

Temperature

• Measured in degrees Fahrenheit (°F), Celsius (°C), or Kelvin (K).

• Conversions:

$K = °C + 273$

$°C = \frac{5}{9}(°F - 32)$

$°F = \frac{9}{5}°C + 32$

Energy and Specific Heat

• Energy: Capacity to do work; measured in joules (J) or calories (cal).

• 1 cal = 4.184 J

• 1 Calorie (Cal, nutritional) = 1000 cal

• Specific Heat (SH): Amount of heat needed to raise the temperature of 1 g of a substance by 1°C.

$\text{Specific heat} = \frac{\text{heat}}{\text{grams} \times \Delta T}$

Table: Specific Heat Values

Substance	Specific Heat (cal/g·°C)
Water (liquid)	1.00
Human body	0.83
Paraffin wax	0.60
Wood, soft	0.34
Wood, hard	0.29
Air	0.24
Aluminum	0.21
Table salt	0.21
Brick	0.20
Stainless steel	0.12
Iron	0.11
Copper	0.092
Silver	0.056
Gold	0.031

States of Matter

• Solid: Definite shape and volume; particles closely packed and fixed.

• Liquid: Adopts shape of container, definite volume; particles loosely packed, random motion.

• Gas: Adopts shape and volume of container; particles far apart, random motion.

Table: Properties of States of Matter

Property	Solid	Liquid	Gas
Shape	Definite	Adopts container	Adopts container
Volume	Definite	Definite	Fills container
Particle arrangement	Closely packed, fixed	Loosely packed, random	Far apart, random
Interparticle forces	Strong	Moderate	Weak

Measuring Matter in Health and Medicine

Accuracy and Precision

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

• Best practice: Take several measurements and average them.

Common Unit Conversions in Health

Unit	SI Equivalent	U.S. Customary Equivalent
Pound (lb)	2.205 lb = 1 kg	1 lb = 16 oz
Quart (qt)	1.057 qt = 1 L	1 qt = 4 cups
Fluid ounce (fl oz)	1 fl oz = 29.6 mL	1 cup = 8 fl oz
Teaspoon (tsp)	1 tsp = 4.93 mL	1 fl oz = 6 tsp
Mile (mi)	1 mi = 1.6 km	1 mi = 5280 ft
Inch (in.)	1 in. = 2.54 cm	1 ft = 12 in.

Calculating Dosages

1. Determine the units for the final answer.

2. Identify the given information.

3. Choose conversion factors to cancel unwanted units.

4. Set up the equation so only the desired unit remains.

Medical Units and Dosing

• Medications may be measured in drops per milliliter (gtt/mL).

• Drop factor depends on IV tubing diameter.

• Active ingredient: Binders may be added to increase pill size.

• Children often receive a percentage of the adult dose based on weight.

• Nutrition labels show % Daily Value (%DV) for nutrients.

Additional info: Some context and examples were inferred and expanded for clarity and completeness, as per academic best practices.