Chemistry - Chapter 1 Notes

Atoms and Molecules

• Properties of matter are determined by the properties of molecules and atoms.
• Understanding matter at the molecular level gives unprecedented control over that matter.
• Atoms are the submicroscopic particles that constitute the fundamental building blocks of ordinary matter.
• Free atoms are rare; they bind together in specific geometrical arrangements to form molecules.
• Liquid water is composed of water molecules, each containing two hydrogen atoms and one oxygen atom held together by a chemical bond.
• Chemistry seeks to understand the behavior of matter by studying the behavior of atoms and molecules.
• Small differences in atoms and molecules can result in large differences in the substances they compose.
• Graphite and diamond are both made of carbon, but their atoms are arranged differently: graphite in sheets and diamond in a three-dimensional structure.

The Scientific Approach to Knowledge

• The scientific approach to knowledge is empirical, based on observation and experiment.
• The scientific method is a process for understanding nature by observing it, studying its behavior, and conducting experiments to test ideas.
• Key characteristics include observation, hypothesis formulation, experimentation, and formulation of laws and theories.
• Observations are descriptions about the characteristics or behavior of nature (also known as data).
• Antoine Lavoisier (1743–1794) noticed no change in the total mass of material within a container during combustion.

Hypothesis

• A hypothesis is a tentative interpretation or explanation of observations.
• Lavoisier hypothesized that during combustion, a substance combines with a component of air.
• A good hypothesis is falsifiable.
• The results of an experiment may support a hypothesis or prove it wrong, requiring the scientist to modify or discard it.

Scientific Law

• A scientific law is a brief statement that summarizes past observations and predicts future ones.
• Law of conservation of mass: “In a chemical reaction, matter is neither created nor destroyed.”
• Scientific laws allow the prediction of future observations and can be tested with experiments.
• Unlike civil or governmental laws, one cannot choose to violate a scientific law.

Theory

• One or more well-established hypotheses may form the basis for a scientific theory.
• A scientific theory is a model for the way nature is and tries to explain not merely what nature does, but why.
• Theories are validated by experiments but can never be conclusively proven because new observations or experiments may reveal a flaw.
• Theories provide general explanations for the characteristics and behavior of nature.
• Dalton’s atomic theory is a model of nature.
• Theories can be used to predict future observations and be tested by experiments.

Conceptual Connection 1.1

• A theory describes what nature does, while a law describes why nature does it.
• A law summarizes a series of related observations, and a theory gives the underlying reasons for them.

The Classification of Matter

• Matter is anything that occupies space and has mass; examples include textbooks, desks, chairs, and bodies.
• Matter can be classified according to its state (physical form) and composition (basic components).
• The states of matter are solid, liquid, and gas.
• The state of matter changes from solid to liquid to gas with increasing temperature.
• Atoms or molecules have different structures in solids, liquids, and gases, leading to different properties.

Solid Matter

• In solid matter, atoms or molecules pack close to each other in fixed locations.
• Atoms and molecules in a solid vibrate but do not move around or past each other.
• A solid has a fixed volume and rigid shape; examples include ice, aluminum, and diamond.
• Solid matter may be crystalline, with atoms or molecules arranged in patterns with long-range, repeating order (e.g., table salt and diamond).
• Amorphous solids do not have long-range order (e.g., glass and plastic).

Liquid Matter

• In liquid matter, atoms or molecules pack about as closely as in solid matter but are free to move relative to each other.
• Liquids have a fixed volume but not a fixed shape.
• A liquid’s ability to flow allows it to assume the shape of its container; examples include water, alcohol, and gasoline at room temperature.

Gaseous Matter

• In gaseous matter, atoms or molecules have a lot of space between them and are free to move relative to one another.
• Gases are compressible.

Classification of Matter by Components

• Matter can be classified as elements, compounds, and mixtures.
• The first division is between a pure substance and a mixture.
• A pure substance is made up of only one component, and its composition is invariant.
• A mixture is composed of two or more components in proportions that can vary from one sample to another.

Classification of Pure Substances

• There are two types of pure substances: elements and compounds.
• This categorization depends on whether or not the substances can be broken down into simpler substances.
• An element is a substance that cannot be chemically broken down into simpler substances.
• Elements are the basic building blocks of matter and are composed of a single type of atom, like helium.
• A compound is a substance composed of two or more elements in fixed, definite proportions.
• Most elements are chemically reactive and combine with other elements to form compounds, such as water and sugar.

Classification of Mixtures

• Mixtures can be categorized into two types: heterogeneous and homogeneous.
• This categorization depends on how uniformly the substances within them mix.

Heterogeneous Mixture

• A heterogeneous mixture is one in which the composition varies from one region of the mixture to another.
• Heterogeneous mixtures are made of multiple substances whose presence can be seen.
• Example: a salt and sand mixture.
• Portions of a sample of a heterogeneous mixture have different compositions and properties.

Homogeneous Mixture

• A homogeneous mixture is made of multiple substances but appears to be one substance.
• All portions of a sample have the same composition and properties (like sweetened tea).
• Homogeneous mixtures have uniform compositions because the atoms or molecules that compose them mix uniformly.

Separating Mixtures

• Mixtures are separable because the different components have different physical or chemical properties.
• Various techniques that exploit these differences are used to achieve separation.
• A mixture of sand and water can be separated by decanting—carefully pouring off the water into another container.

Distillation

• A homogeneous mixture of liquids can usually be separated by distillation, a process in which the mixture is heated to boil off the more volatile (easily vaporizable) liquid.
• The volatile liquid is then recondensed in a condenser and collected in a separate flask.

Filtration

• A mixture of an insoluble solid and a liquid can be separated by filtration—a process in which the mixture is poured through filter paper in a funnel.

Physical and Chemical Changes

Physical Change

• Physical changes alter only the state or appearance of a substance, but not its composition.
• The atoms or molecules that compose a substance do not change their identity during a physical change.
• When water boils, it changes its state from a liquid to a gas, which is a physical change.
• The gas remains composed of water molecules, so this is a physical change.

Chemical Change

• Chemical changes alter the composition of matter.
• During a chemical change, atoms rearrange, transforming the original substances into different substances.
• Rusting of iron is a chemical change.

Physical and Chemical Properties

• A physical property is a property that a substance displays without changing its composition.
• Examples of physical properties include odor, taste, color, appearance, melting point, boiling point, and density.
• The smell of gasoline is a physical property.
• A chemical property is a property that a substance displays only by changing its composition via a chemical change (or chemical reaction).
• The flammability of gasoline is a chemical property.
• Chemical properties include corrosiveness, acidity, and toxicity.

Energy: A Fundamental Part of Physical and Chemical Change

• Energy is the capacity to do work.
• Work is defined as the action of a force through a distance.
• When you push a box across the floor or pedal your bicycle across the street, you have done work.

Types of Energy

• Kinetic energy is the energy associated with the motion of an object.
• Potential energy is the energy associated with the position or composition of an object.
• Thermal energy is the energy associated with the temperature of an object.
• Thermal energy is a type of kinetic energy because it arises from the motion of the individual atoms or molecules that make up an object.

Summarizing Energy

• Energy is always conserved in a physical or chemical change; it is neither created nor destroyed (law of conservation of energy).
• Systems with high potential energy tend to change in a direction that lowers their potential energy, releasing energy into the surroundings.

The Units of Measurement

• In chemistry, units—standard quantities used to specify measurements—are critical.
• The two most common unit systems are the metric system (used in most of the world) and the English system (used in the United States).
• Scientists use the International System of Units (SI), which is based on the metric system.
• The abbreviation SI comes from the French phrase Système International d’ Unités.

Table 1.1 The Standard Units

• 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)

The Meter: A Measure of Length

• The meter (m) is slightly longer than a yard (1 yard is 36 inches, while 1 meter is 39.37 inches).
• 1 meter = \frac{1}{10,000,000} of the distance from the equator to the North Pole (through Paris).
• The International Bureau of Weights and Measures now defines the meter as the distance light travels through a vacuum in \frac{1}{299,792,458} second.

The Kilogram: A Measure of Mass

• The mass of an object is a measure of the quantity of matter within it.
• The SI unit of mass = kilogram (kg)
  • 1 kg = 2.205 lb
• A second common unit of mass is the gram (g).
  • One gram is \frac{1}{1000} kg.
• The weight of an object is a measure of the gravitational pull on its matter.

The Second: A Measure of Time

• Measure of the duration of an event
• SI unit = second (s)
• 1 s is defined precisely as the duration of 9,192,631,770 periods of radiation emitted from a certain transition in a cesium-133 atom.

The Kelvin: A Measure of Temperature

• The kelvin (K) is the SI unit of temperature.
• The temperature is a measure of the average amount of kinetic energy of the atoms or molecules that compose the matter.
• Temperature also determines the direction of thermal energy transfer, or what we commonly call heat.
• Thermal energy transfers from hot to cold objects.
• Kelvin scale (absolute scale) assigns 0 K (absolute zero) to the coldest temperature possible.
• Absolute zero = -273.15°C or -459.67°F is the temperature at which molecular motion virtually stops.

Temperature Scale Equations

• Temperature scale conversion is done with these formulas:
  • °F = 1.8 (°C) + 32
  • K = °C + 273.15

Prefix Multipliers

• The International System of Units uses prefix multipliers with the standard units.
• These multipliers change the value of the unit by powers of 10.
• For example, the kilometer has the prefix kilo meaning 1000 or 10^3.

Table 1.2 SI Prefix Multipliers

• Prefixes and their symbols and multipliers:
  • exa (E): 10^{18}
  • peta (P): 10^{15}
  • tera (T): 10^{12}
  • giga (G): 10^9
  • mega (M): 10^6
  • kilo (k): 10^3
  • deci (d): 10^{-1}
  • centi (c): 10^{-2}
  • milli (m): 10^{-3}
  • micro (\mu): 10^{-6}
  • nano (n): 10^{-9}
  • pico (p): 10^{-12}
  • femto (f): 10^{-15}
  • atto (a): 10^{-18}

Derived Units: Volume and Density

• A derived unit is a combination of other units.
• Volume is a measure of space; it has units of length cubed (cm^3) or liters (L).
• Density is the ratio of a substance’s mass to volume; it has units of mass/volume.
• Density affects if a substance will sink or float in another.
• The less dense substance floats.

Table 1.4 The Density of Some Common Substances at 20 °C

• Substances and their densities (g/cm3):
  • Charcoal (from oak): 0.57
  • Ethanol: 0.789
  • Ice (at 0 °C): 0.917
  • Water (at 4 °C): 1.00
  • Sugar (sucrose): 1.58
  • Table salt (sodium chloride): 2.16
  • Glass: 2.6
  • Aluminum: 2.70
  • Titanium: 4.51
  • Iron: 7.86
  • Copper: 8.96
  • Lead: 11.4
  • Mercury: 13.55
  • Gold: 19.3
  • Platinum: 21.4

Intensive and Extensive Properties

• An intensive property is a characteristic that is independent of the amount of substance.
  • Density is an intensive property.
• An extensive property is a characteristic that is dependent on the amount of substance.
  • Mass is an extensive property.

The Reliability of a Measurement: Significant Figures

• Scientific measurements are reported so that every digit is certain except the last, which is estimated.
• Significant figures deal with writing numbers to reflect precision.
• The precision of a measurement depends on the instrument used to make the measurement.
• The preservation of this precision during calculations can be accomplished by using significant figures.

Counting Significant Figures

• The greater the number of significant figures, the greater the certainty of the measurement.

Significant Figure Rules

1. All nonzero digits are significant.
2. Interior zeroes (zeroes between two nonzero digits) are significant.
3. Leading zeroes (zeroes to the left of the first nonzero digit) are not significant; they only serve to locate the decimal point.
4. Trailing zeroes are categorized as follows:
   • Trailing zeroes after a decimal point are always significant.
   • Trailing zeroes before a decimal point (and after a nonzero number) are always significant.
   • Trailing zeroes before an implied decimal point are ambiguous and should be avoided by using scientific notation.
   • Decimal points are placed after one or more trailing zeroes if the zeroes are to be considered significant.

Example:

• 1200 ambiguous
• 1.2 \times 10^3: 2 significant figures
• 1.20 \times 10^3: 3 significant figures
• 1.200 \times 10^3: 4 significant figures
• 1200.: 4 significant figures

Exact Numbers

• Exact numbers have an unlimited number of significant figures.
  • Accurate counting of discrete objects
  • Defined quantities
  • Integral numbers that are part of an equation
• Some conversion factors are defined quantities, while others are not.

Significant Figures in Calculations

• In calculations using measured quantities, the results of the calculation must reflect the precision of the measured quantities.
• We should not lose or gain precision during mathematical operations.

Rules for Multiplication and Division

• In multiplication or division, the result carries the same number of significant figures as the factor with the fewest significant figures.

Rules for Addition and Subtraction

• In addition or subtraction, the result carries the same number of decimal places as the quantity with the fewest decimal places.
• It is helpful to draw a line next to the number with the fewest decimal places in the answer to determine the number of decimal places in the answer.

Rules for Rounding

• When rounding to the correct number of significant figures:
  • If the last (or leftmost) digit dropped is four or less, round down.
  • If the last (or leftmost) digit dropped is five or more, round up.

Rounding in Multistep Calculations

• To avoid rounding errors in multistep calculations, round only the final answer.
• Keep track of significant figures by underlining the least significant digit in intermediate answers.
• Do not round intermediate steps. If you write down intermediate answers,

Precision and Accuracy

• Accuracy refers to how close the measured value is to the actual value.
• Precision refers to how close a series of measurements are to one another or how reproducible they are.
• Measurements are said to be:
  • precise if they are consistent with one another.
  • accurate only if they are close to the actual value.
• Random error is an error that has the equal probability of being too high or too low. Inaccurate.
• Systematic error is an error that tends toward being either too high or too low. Results are precise but inaccurate.

Solving Chemical Problems

• Many of the problems you will solve in this course are unit conversion problems.
• Using units as a guide to solving problems is called dimensional analysis.
• Units should always be included in calculations; they are multiplied, divided, and canceled like any other algebraic quantity.

Dimensional Analysis

• A unit equation is a statement of two equivalent quantities, such as 2.54 cm = 1 in.
• A conversion factor is a fractional quantity of a unit equation with the units we are converting from on the bottom and the units we are converting to on the top.
• Most unit conversion problems take the following form:
  • Information given × conversion factor(s) = information sought
  • Given unit × \frac{desired unit}{given unit} = desired unit

General Problem Solving Strategy

• Identify the starting point, the given information.
• Identify the end point, what you must find.
• Devise a way to get from the starting point to the end point using what is given as well as what you already know or can look up, the conceptual plan.
• The general steps are sort information, strategize via conceptual plan, solve by following your conceptual plan, and then check that your answer makes sense.

Units Raised to a Power

• When building conversion factors for units raised to a power, remember to raise both the number and the unit to the power.
• For example, to convert from in2 to cm2, we construct the conversion factor as follows:

\frac{(2.54 cm)^2}{(1 in)^2} = \frac{(2.54 cm)}{(1 in)} \times \frac{(2.54 cm)}{(1 in)} = \frac{6.45 cm^2}{1 in^2}

Problems Involving Equations

• Solve the same way as problems with conversion factors
• Same four steps: sort, strategize, solve, check
• Equation takes us from the given quantities to the find quantity
• Use the equation as the conversion, filling in the given quantities for given and solve for the find

Analyzing Data

• Learning to analyze and interpret data is an important scientific skill.
• Example: Understanding the composition of water.
• Suppose you are an early chemist trying to understand the composition of water.
• Water samples are decomposed into hydrogen and oxygen.

Interpreting Data Pattern Recognition

• The sum of the masses of oxygen and hydrogen always sum to the mass of the water sample.
• The ratio of masses of oxygen to hydrogen is the same for each sample with small variations due to experimental error. The ratio is ~8.

Interpreting Graphs

• Data are often visualized using graphs or images.
• Scientists must constantly analyze and interpret graphs.
• Carbon dioxide is a greenhouse gas that has been rising as result of the burning of fossil fuels (such as gasoline and coal).
• First examine the x and y axes to understand what each represents.
• Also examine the numerical range of the axes.
• Note if the y axis start at a non-zero value.
• The increase in carbon dioxide has not been constant over time.
• The rate of increase — slope of the line — has intensified since about 1960.