General Chemistry Fundamentals: Measurement, Atomic Structure, and the Periodic Table

Chemistry Fundamentals

Introduction

This study guide covers foundational topics in General Chemistry, including measurement, lab equipment, significant figures, density, scientific notation, dimensional analysis, atomic structure, isotopes, and the periodic table. Mastery of these concepts is essential for success in college-level chemistry courses.

Measurement in Chemistry

Accuracy vs. Precision

Understanding the difference between accuracy and precision is crucial for interpreting experimental data.

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

• Precision: How repeatable measurements are, regardless of their closeness to the true value.

• Data: Precision is related to the repeatability of experimental data.

• Instruments: Precision can also refer to the number of decimal places an instrument can record.

Example: Measuring the boiling point of water at 98.5°C, 101.0°C, 100.2°C, and 100.0°C: These values are precise (close to each other) but not necessarily accurate (close to the true value of 100.0°C).

Lab Equipment

Proper use of laboratory equipment is essential for obtaining accurate and precise measurements.

• Estimated digit: The last digit in a measurement, which is an estimate between the smallest graduations.

• Calibration: The process of configuring an instrument to provide results within an acceptable range.

• Common glassware: Beaker, Erlenmeyer flask, graduated cylinder, volumetric flask, burette, pipette.

Glassware	Name	Use
	Beaker	Holding and mixing liquids
	Erlenmeyer Flask	Mixing and heating liquids
	Graduated Cylinder	Measuring liquid volume accurately
	Volumetric Flask	Preparing solutions of precise volume
	Burette	Dispensing precise volumes, especially in titrations
	Pipette	Transferring small, precise volumes of liquid

Reading Measurements

• Always record all certain digits plus one estimated digit.

• Precision depends on the instrument used (e.g., a burette is more precise than a beaker).

Example: Comparing rulers with different scales to determine which provides the most precise measurement.

Density

Density is a fundamental property used to identify substances and solve quantitative problems.

• Definition: Density is the mass of a substance per unit volume.

• Formula:

$D = \frac{\text{Mass}}{\text{Volume}}$

• Units: g/mL or g/cm3 for liquids and solids; g/L for gases.

• Density is an intensive property (independent of sample size).

Example: Calculate the volume occupied by 35.2 g of a substance with a density of 1.60 g/mL.

$\text{Volume} = \frac{35.2\ \text{g}}{1.60\ \text{g/mL}} = 22.0\ \text{mL}$

Substance	Density (g/mL) at 25°C
Mercury	13.6
Water	1.00
Hexane (non-polar)	0.660
Ethyl Alcohol (polar)	0.789
Dichloromethane (polar)	1.33
Aluminum	2.70
Bromine (non-polar)	2.929
Gold	19.3

Application: Identifying unknown liquids by comparing their densities to known values.

Scientific Notation

Scientific notation is used to express very large or very small numbers efficiently.

• Format: $a \times 10^n$, where $1 \leq a < 10$ and $n$ is an integer.

• Positive exponent: Number is greater than 1.

• Negative exponent: Number is less than 1.

Example: 4,500 = $4.5 \times 10^3$; 0.0677 = $6.77 \times 10^{-2}$

Standard	Scientific
43,900	$4.39 \times 10^4$
0.00943	$9.43 \times 10^{-3}$
2.62 × 103	2,620
8.65 × 104	86,500

Significant Figures

Significant figures (sig figs) indicate the precision of a measured or calculated quantity.

• Rules for counting sig figs:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

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

• Exact numbers (e.g., counted items, defined conversions) have infinite sig figs.

Rounding: The answer should have the same number of decimal places (for addition/subtraction) or significant figures (for multiplication/division) as the least precise measurement.

Example:

Problem	What is the least precise place value?	Full Calculated Answer	Answer rounded to correct # of sig figs
12.4 cm + 7.89 cm	1 decimal place	20.29 cm	20.3 cm
210 cm - 12.3 cm	no decimal places	197.7 cm	198 cm

Dimensional Analysis

Dimensional analysis is a systematic method for converting between units using conversion factors.

• Always start with what you are given.

• Choose appropriate conversion factors.

• Ensure units cancel diagonally.

• Continue until the desired unit is reached.

Metric System Prefixes: Common prefixes include kilo- (k, $10^3$), centi- (c, $10^{-2}$), milli- (m, $10^{-3}$), micro- ($\mu$, $10^{-6}$), nano- (n, $10^{-9}$), pico- (p, $10^{-12}$).

Unit Symbol	Unit Conversion	Metric Prefix	Factor
m	1 m = 100 cm	kilo- (k)	$10^3$
L	1 L = 1000 mL	centi- (c)	$10^{-2}$
g	1 kg = 1000 g	milli- (m)	$10^{-3}$
s	1 min = 60 s	micro- ($\mu$)	$10^{-6}$

Example: Convert 24 kilometers per minute to meters per second.

$24\ \text{km/min} \times \frac{1000\ \text{m}}{1\ \text{km}} \times \frac{1\ \text{min}}{60\ \text{s}} = 400\ \text{m/s}$

Atomic Structure

Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons.

Particle	Symbol	Charge	Location
Proton	p+	+1	Nucleus
Neutron	n0	0	Nucleus
Electron	e-	-1	Electron cloud

• Protons and neutrons have similar mass; electrons are much lighter.

• The number of protons defines the atomic number (Z) and the identity of the element.

• The sum of protons and neutrons gives the mass number (A).

• Neutral atoms have equal numbers of protons and electrons.

• Ions are formed by gaining or losing electrons.

Example: The nuclear symbol for an atom is written as:

$^{A}_{Z}\text{X}^{\text{charge}}$

Isotopes

Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons.

• Isotopes have the same atomic number but different mass numbers.

Isotopic Notation	Protons	Neutrons	Electrons
$^{24}_{12}\text{Mg}$	12	12	12
$^{36}_{17}\text{Cl}^-$	17	19	18
$^{19}_{9}\text{F}$	9	10	9

The Periodic Table

Metals, Nonmetals, and Metalloids

• Metals: Good conductors, malleable, ductile, shiny, mostly solids at room temperature.

• Nonmetals: Poor conductors, brittle, dull, can be solids, liquids, or gases.

• Metalloids: Properties intermediate between metals and nonmetals; found along the staircase line (e.g., Si, B, As).

Periods vs. Groups

• Periods: Horizontal rows; elements in the same period have the same number of electron shells.

• Groups (families): Vertical columns; elements in the same group have similar chemical properties and the same number of valence electrons.

Element	Group	# Valence Electrons	Charge in Ionic Compounds	Cation or Anion?
Rb	1	1	+1	Cation
N	15	5	-3	Anion
Ne	18	8	0	Neither
Mg	2	2	+2	Cation
Fe	8	varies	+2, +3	Cation
F	17	7	-1	Anion

Lewis Dot Structures

Lewis dot structures represent the valence electrons of atoms and predict the ions they form.

Element Symbol	Lewis Model of Neutral Atom	Lewis Model of Ion That Will Form
Ca	Ca with 2 dots	Ca2+ (no dots)
P	P with 5 dots	P3- (8 dots)
N	N with 5 dots	N3- (8 dots)

Example: The Lewis structure for Cl- shows 8 valence electrons around the symbol Cl.

Additional info: Some tables and diagrams were inferred and expanded for clarity. This guide covers all major introductory topics in General Chemistry as presented in the provided materials.