Chapter E and 1 quiz review

Chapter E Essentials: Units, Measurement, and Problem Solving

Units of Measurement

• Importance of units in chemistry as standard quantities for specifying measurements.

• Two common unit systems:

  • Metric system: Used globally.

  • English system: Used primarily in the U.S.

Table E.1: SI Base Units

• Length: Meter (m)

• Mass: Kilogram (kg)

• Time: Second (s)

• Temperature: Kelvin (K)

• Amount of Substance: Mole (mol)

• Electric Current: Ampere (A)

• Luminous Intensity: Candela (cd)

important conversions to know for quiz

• 1 L = 1000mL

• 1 km= 1000g

• 1m= 100 cm

The Meter: A Measure of Length

• The meter (m) is 39.37 inches, slightly longer than a yard (36 inches).

• Important conversions:

  • 1 m = 100 cm

  • 1 m = 1000 mm

  • 1 km = 1000 m

The Kilogram: A Measure of Mass

• Mass measures the quantity of matter in an object.

• SI unit: Kilogram (kg)

  • Equivalence: 1 kg = 2 lb 3 oz

• Alternative mass unit: Gram (g)

  • 1 g = 1/1000 kg

• Weight measures gravitational pull on mass.

The Second: A Measure of Time

• Represents the duration of an event.

• SI unit: Second (s)

Volume

• The space a substance occupies; SI unit is m³, commonly expressed in liters (L), milliliters (mL), and cm³.

The Kelvin: A Measure of Temperature

• Kelvin (K) is the SI unit of temperature.

• Equation to find celcius: C= (Farenheit-32)/1.8

• equation to find kelvin: K= celsius -273.15

• Measures average kinetic energy of atoms/molecules.

• Thermal energy flows from hotter to colder objects.

Absolute Temperature Scale

• 0 K is absolute zero, the coldest temperature where molecular motion virtually stops.

Temperature Scales

• Fahrenheit (°F): Degree size is five-ninths that of Celsius (°C).

• Celsius (°C) and Kelvin (K) are equivalent in degree size.

• Temperature conversion methods outlined.

Precision and Accuracy

• Accuracy: Closeness of measured value to actual value.

• Precision: Closeness of repeated measurements.

Counting Significant Figures

• Significant figures reflect the precision of a measurement.

• Instrument precision affects measurements and subsequent calculations.

• Guidelines for significant figures include:

  • Nonzero digits are always significant.

  • Interior zeros are significant.

  • Leading zeros are not significant.

  • Trailing zeros may or may not be significant based on context.

Exact Numbers

• Exact numbers have an unlimited number of significant figures, such as counts or defined quantities.

Rules for Calculations

• Multiplication/Division Rule: Result carries the same significant figures as the factor with the fewest.

• Addition/Subtraction Rule: Result carries the same decimal places as the least precise quantity.

Rules for Rounding

• Round down if the last digit dropped is 4 or less; round up if 5 or more.

Rounding in Multistep Calculations

• the last order of operation determines which rule applies in terms of significant figures.

Energy

• Energy: Capacity to do work—depends on force applied over distance.

• Kinetic Energy: Related to the motion of an object.

• Potential Energy: Related to position or composition of an object. also known is resting state.

• Thermal Energy: A type of kinetic energy from motion of atoms/molecules.

Summarizing Energy

• The law of conservation of energy states energy is never created or destroyed.

• Systems with high potential energy will change to lower energy states, releasing energy.

Energy Conversion Factors

Table E.5

• 1 calorie (Cal) = 4.184 joules (J)

• 1 Calorie (Cal) or kilocalorie (kcal) = 1000 cal = 4184 J

• 1 kilowatt-hour (kWh) = 3.60 x 10^6 J

Chapter 1

Properties of Matter

• The properties of matter are fundamentally determined by the properties of molecules and atoms.

• Atoms and molecules influence how matter behaves, making them the building blocks of ordinary matter.

• Free atoms are rare in nature; they typically bind together in specific geometrical arrangements to form molecules.

Atoms and Elements

• An atom is the smallest identifiable unit of an element.

• There are approximately 91 different naturally occurring elements and over 20 synthetic elements, which are not found in nature.

Classification of Matter

• Matter is defined as anything that occupies space and possesses mass.

• It can be classified based on its state (solid, liquid, gas) and its composition.

  • States of Matter:

    • Solid, Liquid, Gas

  • State Classifications:

    • Solid: Atoms or molecules tightly packed in fixed locations.

    • Liquid: Atoms or molecules close together but free to move relative to one another.

    • Gas: Atoms or molecules far apart and free to move, making them compressible.

Solid Matter

• Characteristics of solid matter:

  • Atoms or molecules are fixed near each other with very little kinetic energy, resulting in a fixed volume and rigid shape.

  • Examples include ice, aluminum, and diamond.

• Solid matter can be:

  • Crystalline: Atoms or molecules organized in a repeating pattern (e.g., table salt, diamond).

  • Amorphous: Lacking long-range order (e.g., glass and plastic).

Gaseous Matter

• Gases consist of atoms or molecules with significant space between them, allowing them to move freely relative to one another, making them compressible.

Classification of Matter by Components

• Matter can also be classified based on its composition into elements, compounds, and mixtures.

• Pure Substance vs Mixture:

  • A pure substance has a uniform composition and cannot be separated into different components by physical means.

  • A mixture is composed of two or more substances that retain their individual properties and can vary in composition.

Classification of Pure Substances

• Element: A fundamental substance that cannot be chemically broken down into simpler substances.

  • Composed of a single type of atom (e.g., helium).

• Compound: A substance formed from two or more elements in fixed proportions.

  • Most elements are reactive and combine to form compounds (e.g., water, sugar).

Classification of Mixtures

• Mixtures are categorized into:

  • Heterogeneous Mixtures: visually distinguishable, like fruit salad

  • Homogeneous Mixtures: Composition is uniform throughout (e.g., salt water mixed).

The Scientific Approach to Knowledge

• Key characteristics of the scientific method:

  • Observation

  • Formulation of hypotheses

  • Experimentation

  • Formulation of laws and theories.

Modern Atomic Theory and Laws

• The atomic theory, stating that all matter is composed of atoms, evolved from observations and laws:

  • Laws guiding this theory:

    • The law of conservation of mass

    • The law of definite proportions

    • The law of multiple proportions

The Law of Conservation of Mass

• stating matter is neither created nor destroyed in chemical reactions, ensuring total mass remains constant during reactions.

The Law of Definite Proportions

• indicates all samples of a specific compound have the same proportions of elements by mass

  • water will always have 2 H and 1 O

The Law of Multiple Proportions

• when two elements form multiple compounds, the masses of one element that combine with a fixed mass of the other can be expressed as a ratio of small whole numbers (e.g., carbon monoxide vs carbon dioxide).

The Neutrons

• Neutrons have a mass similar to protons but no charge, affecting atomic mass.

• only protons and neutrons affect atomic mass. electrons do not.

Subatomic Particles

• Atoms consist of three primary subatomic particles:

  • Protons

  • Neutrons

  • Electrons

Elements and Their Protons

• The identity of an atom is defined by the number of protons in its nucleus, known as its atomic number (Z).

• atomic number = the number of protons in nucleus

Isotopes and Mass Number

• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

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

Ions: Gaining or Losing Electrons

• Neutral atoms have an equal number of protons and electrons; losing or gaining electrons results in ions.

  • Cations: Positively charged ions (e.g., Na+).

  • Anions: Negatively charged ions (e.g., F–).