Chemistry Basics: Matter and Measurement

Effective Study Strategies for Chemistry

Success in chemistry requires active engagement with the material, including reading, note-taking, and problem-solving. Passive techniques, such as re-reading or re-listening, are less effective than active learning.

• Definitions: Always learn bold-faced terms in the textbook.

• Concepts: Explain concepts to others to test understanding.

• Problems: Practice problems from class and assignments.

• Pattern Recognition: Chemistry involves recognizing patterns in problems, molecular structures, and reactions.

• Active Learning: Study daily, review previous material, and avoid rote memorization.

Zones of Learning

Learning occurs outside the comfort zone, requiring students to cross the fear zone and enter the learning or growth zone.

• Comfort Zone: Familiar, but not conducive to growth.

• Learning Zone: Where new material is mastered.

• Growth Zone: Advanced mastery and application.

Classifying Matter

Pure Substances and Mixtures

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

• Pure Substance: Made of only one type of substance. Includes elements and compounds.

• Element: Simplest type of matter, made of one type of atom.

• Compound: Pure substance made of two or more elements chemically joined.

• Mixture: Combination of two or more substances that can be separated.

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

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

Elements, Compounds, and the Periodic Table

Periodic Table Structure

The periodic table organizes all known elements by their properties.

• Group: Vertical column; elements in a group have similar chemical behaviors.

• Period: Horizontal row; numbered 1 to 7.

• Staircase Line: Separates metals from nonmetals; elements bordering the line (except Al) are metalloids.

• Element Symbols: Derived from English or Latin names (e.g., Na for sodium).

Chemical Formulas

Elements combine in fixed ratios to form compounds, represented by chemical formulas.

• Example: (water), (carbon dioxide)

Physical and Chemical Changes

Types of Changes

Matter can undergo physical or chemical changes.

• Physical Change: Alters form or appearance, but not identity (e.g., melting ice).

• Chemical Change: Alters chemical composition; new substances are formed (e.g., rusting iron).

Chemical Equations

Chemical equations describe reactions, showing reactants and products.

• Reactants: Substances that undergo change.

• Products: Substances formed.

• Arrow (): Indicates "react to form".

• Physical States: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

Balancing Chemical Equations

Equations must be balanced to obey the law of conservation of mass.

• Law of Conservation of Mass: Matter is neither created nor destroyed.

• Coefficients: Numbers placed in front of formulas to balance atoms.

Example:

Measurement and Units

SI Units and Metric Prefixes

The SI system is used for scientific measurements.

• Mass: Kilogram (kg)

• Volume: Liter (L)

• Length: Meter (m)

Dimensional Analysis and Conversion Factors

Conversion factors are used to change units.

• Equivalencies:

• Dimensional Analysis: Multiply by conversion factors to cancel units.

Significant Figures

Significant figures reflect the certainty of measurements.

• All nonzero digits are significant.

• Zeros: Significance depends on position (see table below).

• Exact Numbers: Have infinite significant figures (e.g., counting objects).

Scientific Notation

Used to express very large or small numbers.

• General Form:

• Positive Exponent: Number is large.

• Negative Exponent: Number is small.

Percent Calculations

Percent means "per hundred" and is used to compare quantities.

• Fraction to Percent: Divide numerator by denominator, multiply by 100.

• Decimal to Percent: Multiply by 100.

Matter: Mass, Volume, Density, and Temperature

Mass and Weight

Mass is the amount of matter in an object; weight depends on gravity.

• Unit for Mass: Gram (g)

• Weight: Varies with location due to gravity.

Volume

Volume is the space occupied by matter.

• Lab Unit: Milliliter (mL)

• Clinical Unit: Cubic centimeter (cc or cm3)

• 1 mL = 1 cm3

Density and Specific Gravity

Density is the ratio of mass to volume; specific gravity compares density to water.

• Density Formula:

• Water's Density:

• Specific Gravity:

Temperature Scales

Temperature measures hotness or coldness; common scales are Celsius, Fahrenheit, and Kelvin.

• Normal Body Temperature: or

• Kelvin Conversion:

• Celsius to Fahrenheit:

Energy and Heat

Energy is the capacity to do work; heat is kinetic energy transferred between objects.

• SI Unit: Joule (J)

• Calorie: Energy to raise 1 g of water by 1°C

• Specific Heat: Amount of heat needed to raise 1 g of substance by 1°C

States of Matter

Matter exists as solids, liquids, or gases, each with distinct particle arrangements and properties.

• Solid: Definite shape and volume; particles tightly packed.

• Liquid: Definite volume, takes shape of container; particles move freely.

• Gas: No definite shape or volume; particles far apart and move rapidly.

Measuring Matter: Accuracy, Precision, and Units

Accuracy and Precision

Accuracy is closeness to true value; precision is consistency among measurements.

• Best Practice: Take multiple measurements and average values.

Common Unit Conversions

Dosage Calculations

Health professionals use SI and U.S. units; IV medications are measured in drops per mL (gtt/mL).

• Drop Factor: Used to calculate drip rates.

• Dosage Calculation Steps:

  1. Determine units for final answer.

  2. Identify given information.

  3. Find conversion factors to cancel units.

  4. Set up equation so only desired units remain.

Atoms and Their Components

Subatomic Particles

Atoms are composed of protons, neutrons, and electrons.

• Proton: Positive charge (+1), mass ≈ 1 amu

• Neutron: No charge (0), mass ≈ 1 amu

• Electron: Negative charge (-1), mass ≈ 1/2000 amu

• Atoms are neutral: Number of protons = number of electrons

Atomic Structure

Protons and neutrons are in the nucleus; electrons occupy the electron cloud.

• Nucleus: Dense center of atom

• Electron Cloud: Region where electrons are found

• Atomic Mass Unit (amu): 1 amu = 1/12 mass of carbon atom

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons; defines element

• Mass Number (A): Number of protons + neutrons

• Symbolic Notation:

Isotopes and Atomic Mass

Isotopes are atoms of the same element with different numbers of neutrons.

• Isotope Notation: Carbon-12, Carbon-13

• Atomic Mass: Weighted average of all naturally occurring isotopes

Radioactivity and Radioisotopes

Radioactivity

Radioactivity is energy released from unstable atomic nuclei.

• Radioisotopes: Isotopes that emit radiation

• Forms of Radiation:

  • Alpha (): Positively charged

  • Beta (): Negatively charged

  • Gamma (): Neutral

  • Positron: Positive charge

  • Neutron: No charge

Electron Arrangements and the Octet Rule

Energy Levels and Electron Configuration

Electrons occupy energy levels; stability is achieved by filling these levels.

• Energy Levels: Electrons fill lowest energy levels first

• Maximum Electrons per Level:

• Valence Level: Highest energy level with electrons

• Valence Electrons: Electrons in outermost level; involved in bonding

Octet Rule

Atoms tend to achieve eight electrons in their valence shell for stability.

• Noble Gases: Group 18; have eight valence electrons (except He)

• Octet Rule: Atoms gain, lose, or share electrons to achieve eight valence electrons

Ion Formation and Ionic Compounds

Ions

Ions are charged atoms formed by gaining or losing electrons.

• Anion: Gains electrons; negative charge

• Cation: Loses electrons; positive charge

• Isoelectronic: Ions with same electron count as nearest noble gas

Trends in Ion Formation

• Groups 1A, 2A, 3A: Form cations with charge equal to group number

• Groups 5A, 6A, 7A: Form anions with charge equal to 8 minus group number

• Transition Metals: May form multiple ions; charge not predictable from group

• Polyatomic Ions: Groups of atoms with overall charge (e.g., )

Naming Ions

• Metal Ions: Add "ion" to metal name (e.g., sodium ion)

• Transition Metals: Use Roman numeral for charge (e.g., iron(III) ion)

• Nonmetals: Suffix "-ide" (e.g., chloride)

• Polyatomic Ions: Names often end in "-ate" or "-ite"

Ionic Compounds

Ionic compounds form from the transfer of electrons between metals and nonmetals.

• Electrostatic Attraction: Holds cations and anions together

• Formula: Net charge must be zero; use subscripts to indicate number of ions

Naming Ionic Compounds

• Combine names of cation and anion; omit "ion"

• Use Roman numerals for transition metals

• Polyatomic ion names remain unchanged

Covalent Compounds and Bonding

Covalent Bond Formation

Nonmetals form covalent bonds by sharing valence electrons to achieve an octet.

• Bonding Pair: Shared pair of electrons; represented by a dash

• Molecule: Fundamental unit of covalent compound

Lewis Structures

Lewis-dot symbols show valence electrons and bonding.

• Step 1: Draw electron-dot symbols

• Step 2: Line up atoms

• Step 3: Draw bonds

• Step 4: Confirm octets

• Step 5: Check structure

Naming Covalent Compounds

• Binary Compounds: Name first element, then second with "-ide"

• Use Greek prefixes to indicate number of each element

• Traditional Names: Some compounds have common names (e.g., water, ammonia)

The Mole: Counting Atoms and Compounds

The mole is a fundamental unit for counting atoms and molecules in chemistry.

• 1 mole = particles (Avogadro's number)

• Used to relate mass, number of particles, and volume in chemical calculations

Additional info: Some context and examples were inferred to ensure completeness and clarity for GOB Chemistry students.