The Study of Chemistry

The Atomic and Molecular Perspective of Chemistry

Chemistry

• Is the study of matter, its properties, and the changes that matter undergoes.

Chemistry also helps explain matter's qualities in terms of atoms, it's practically infinitesimally small building units.

Matter

• Is the physical material of the universe; it is anything that has mass and occupies space.

Property

• Is any characteristic that allows us to recognize a particular type of matter and to distinguish it from other types.

Elements

• All matter is composed of combinations of only about 100 substances.

Molecules

• Two or more atoms are joined in specific shapes.

Ethanol

• Contains one oxygen atom, depicted by one red sphere.

• Is the alcohol in beverages such as beer and wine.

• Is consumed throughout the world.

Ethylene glycol

• Contains two oxygen atoms.

• Is a viscous liquid used as automobile antifreeze.

• You should never consume ethylene glycol because it is highly toxic.

Chemists must manage the composition or structure of molecules to create new substances with various qualities.

For instance, the widely used painkiller aspirin was first created in 1897 in a successful attempt to outperform a natural substance collected from willow bark.

• As we proceed with our study of chemistry, we will find ourselves thinking in two realms: the macroscopic realm of ordinary-sized objects 1macro = large2 and the submicroscopic realm of atoms and molecules.

Why Study Chemistry?

• Chemistry is crucial to improving health care, conserving natural resources, protecting the environment, and supplying society with energy. Using chemistry, we developed and enhanced medications, fertilizers, insecticides, polymers, solar panels, LEDs, and building materials. We also found compounds that endanger humans and the environment. We must ensure the products we touch are safe.

• Chemistry is central to a fundamental understanding of governing principles in many science-related fields.

Chemistry is central to our understanding of the world around us.

Energy

• Solar panels are composed of specially treated silicon.

Biochemistry

• The flash of the firefly results from a chemical reaction in the insect.

Technology

• OLEDs (organic light-emitting diodes) are used in high-end cell phone, tablet, and television displays.

Medicine

• Chemists are constantly striving to design new and improved drugs for treating disease.

Chemistry and the Chemical Industry

• The chemical industry in the United States is estimated to be an $800 billion enterprise that employs over 800,000 people and accounts for 14% of all U.S. exports.

CHEMISTS

• Chemists work in industry, government, and academia. Industry lab chemists develop new goods, analyze materials, and help customers use them.

• Chemists work in universities and government agencies like the National Institutes of Health, Department of Energy, and Environmental Protection Agency. Teaching, medicine, biomedical research, information science, environmental work, technical sales, government regulatory bodies, and patent law all benefit from chemistry degrees.

Chemists do three things:

• They make new types of matter: materials, substances, or combinations of substances with desired properties.

• Measure the properties of matter

• Develop models that explain and/or predict the properties of matter.

Classifications of Matter

States of Matter

• A sample of matter can be a gas, a liquid, or a solid. These three forms, called the states of matter, differ in some of their observable properties.

• Gas (also known as vapor) has no fixed volume or shape; rather, it uniformly fills its container. A gas can be compressed to occupy a smaller volume, or it can expand

• to occupy a larger one.

• Liquid has a distinct volume independent of its container, assumes the shape of the portion of the container it occupies, and is not compressible to any appreciable extent.

• Solid has both a definite shape and a definite volume and is not compressible to any appreciable extent.

The properties of the states of matter can be understood on the molecular level.

GAS

• The molecules collide frequently and are widely apart.

• Compressing a gas reduces the spacing between molecules and increases their collision frequency without changing their size or form.

LIQUID

• The molecules are close but move fast.

• Liquids spill easily because molecules glide over each other during fast movement.

SOLID

• The molecules are kept closely together, usually in defined patterns so they can only wiggle minimally. Thus, the lengths between molecules are similar in liquid and solid phases, yet molecules are mostly held in place in solids but have considerable freedom of motion in liquids.

Pure Substances

• (usually referred to simply as a substance) is matter that has distinct properties and a composition that does not vary from sample to sample. Water and table salt (sodium chloride) are examples of pure substances.

Elements

• Are substances that cannot be decomposed into simpler substances. On the molecular level, each element is composed of only one kind of atom.

• The 118 known elements vary in quantity. The Milky Way galaxy has 74% hydrogen and 24% helium. Nearer to home, five elements—oxygen, silicon, aluminum, iron, and calcium—make up over 90% of Earth's crust (including seas and atmosphere), and three—oxygen, carbon, and hydrogen—make up over 90% of the human body.

Compounds

• Are substances composed of two or more elements; they contain two or more kinds of atoms. Water, for example, is a compound composed of two elements: hydrogen and oxygen.

• Most elements form compounds. Hydrogen gas burns in oxygen gas to make water. Water can be degraded by electrical current.

Law of Constant Composition (or the law of definite proportions)

• The observation that the elemental composition of a compound is always the same.

• Although this law has been known for 200 years, some people believe that molecules manufactured in the lab and those found in nature are fundamentally different. Pure compounds have the same composition and properties regardless of their origin. Chemists and nature employ the same components and follow the same laws. Two materials with different compositions or qualities either have different compounds or are pure.

French chemist Joseph Louis Proust (1754–1826)

• first stated the law in about 1800.

Mixtures

are combinations of two or more substances in which each substance retains its chemical identity.

• The substances making up a mixture are called components of the mixture.

• Some mixtures do not have the same composition, properties, and appearance throughout.

• Homogeneous mixtures are also called solutions.

• Mixtures that are uniform throughout are homogeneous.

All pure matter is classified ultimately as either an element or a compound.

PROPERTIES OF MATTER

• The properties of matter can be categorized as physical or chemical.

• Physical properties can be observed without changing the identity and composition of the substance. These properties include color, odor, density, melting point, boiling point, and hardness.

Chemical properties

• Describe the way a substance may change, or react, to form other substances. A common chemical property is flammability, the ability of a substance to burn in the presence of oxygen.

• Some properties, such as temperature and melting point, are intensive properties.

Intensive properties

• Do not depend on the amount of sample being examined and are particularly useful in chemistry because many intensive properties can be used to identify substances.

• Extensive properties depend on the amount of sample, with two examples being mass and volume. Extensive properties relate to the amount of substance present.

Physical and Chemical Changes

Physical Change

• a substance changes its physical appearance but not its composition. (That is, it is the same substance before and after the change.)

• The evaporation of water is a physical change.

All changes of state (for example, from liquid to gas or from liquid to solid) are physical changes.

Chemical Change

• (also called a chemical reaction), a substance is transformed into a chemically different substance. When hydrogen burns in air, for example, it undergoes a chemical change because it combines with oxygen to form water.Chemical changes can be dramatic.

Separation of Mixtures

• Using properties, we can separate a mixture. For instance, iron and gold filings can be separated by color. Using a magnet to attract iron filings and leave gold ones would be easier. We can also use a significant chemical difference between these metals: Gold resists most acids. Thus, a suitable acid would dissolve the iron in our mixture, leaving solid gold. Filtration separated them.

Distillation

• An important method of separating the components of a homogeneous mixture.

• A process that depends on the different abilities of substances to form gasses.

Chromatography Technique

1. Mixture to be separated is dissolved in the solvent.

2. Solvent is added throughout the process.

3. Components separate.

4. Each component is collected as it reaches the bottom of the column.

THE NATURE OF ENERGY

• Energy is defined as the capacity to do work or transfer heat.

• Work is the energy transferred when a force exerted on an object causes a displacement of that object, and heat is the energy used to cause the temperature of an object to increase.

• We define work, w, as the product of the force exerted on the object, F,is defined as any push or pull exerted on the object and the distance, d, that it moves: w = Fxd

Kinetic Energy and Potential Energy

• Objects, whether they are automobiles, soccer balls, or molecules, can possess kinetic energy, the energy of motion. The magnitude of kinetic energy, Ek, of an object depends on its mass, m, and velocity, v: Ek = 1/2 mv^2

• Thus, the kinetic energy of an object increases as its velocity or speed increases.

• Example: A car has greater kinetic energy moving at 65 miles per hour (mi/h) than it does at 25 mi/h. For a given velocity, the kinetic energy increases with increasing mass. Thus, a large truck traveling at 65 mi/h has greater kinetic energy than a motorcycle traveling at the same velocity because the truck has the greater mass.

Potential Energy

• All other forms of energy—the energy stored in a stretched spring, in a weight held above your head, or in a chemical bond.

• is in essence, the “stored” energy that arises from the attractions and repulsions an object

• experiences in relation to other objects.

Atoms and molecules hardly ever interact with one another as a result of gravitational forces. When working with atoms and molecules, electrical charge-based forces are more significant.

Electrostatic Potential Energy

• most important forms of potential energy in chemistry.

• which arises from the interactions between charged particles.

• Many substances, fuels for example, release energy when they react. The chemical energy of a fuel is due to the potential energy stored in the arrangements of its atoms.

UNITS OF MEASUREMENT

QUANTITATIVE

• Associated with numbers.

• When a number represents a measured quantity, the units of that quantity must be specified.

Metric System

• The units used for scientific measurements are those of the metric system.

• The metric system, developed in France during the late eighteenth century, is used as the system of measurement in most countries. The United States has traditionally used the English system, although use of the metric system has become more common.

SI Units

• In 1960, an international agreement was reached specifying a particular choice of metric units for use in scientific measurements.

• After the French Système International d’Unités. This system has seven base units from which all other units are derived.

The scientific method

1. Collect information (via observations of natural phenomena and experiments)

2. Formulate a hypothesis (or hypotheses)

3. Test the hypothesis (via experiments)

4. Formulate a theory (based on the most successful hypotheses)

5. Repeatedly test theory (modify as needed to match experimental results, or reject)

Length and Mass

• The SI base unit of length is the meter, a distance slightly longer than a yard. Mass is a measure of the amount of material in an object. The SI base unit of mass is the kilogram (kg), which is equal to about 2.2 pounds (lb). This base unit is unusual because it uses a prefix, kilo-, instead of the word gram alone. We obtain other units for mass by adding prefixes to the word gram.

Temperature

• A measure of the hotness or coldness of an object, is a physical property that determines the direction of heat flow.

• Heat always flows freely from a higher-temperature substance to a lower-temperature substance. Thus, the flood of heat we feel when we touch a hot object indicates that the thing is hotter than our hand.

• Science uses Celsius and Kelvin scales.

Celsius scale

• Was originally based on the assignment of 0 °C to the freezing point of water and 100 °C to its boiling point at sea level.

• Absolute zero has the value -273.15 °C.

Kelvin scale

• Is the SI temperature scale, and the SI unit of temperature is the kelvin (K). Zero on the Kelvin scale is the temperature at which all thermal motion ceases, a temperature referred to as absolute zero.

• The Celsius and Kelvin scales have equal-sized units—that is, a kelvin is the same size as a degree Celsius.

• K = °C + 273.15

Derived SI Units

• The SI base units are used to formulate.

• A derived unit is obtained by multiplication or division of one or more of the base units.

• For example, speed is defined as the ratio of distance traveled to elapsed time. Thus, the derived SI unit for speed is the SI unit for distance (length), m, divided by the SI unit for time, s, which gives m/s, read “meters per second.” Two common derived units in chemistry are those for volume and density.

Volume

• A cube's volume is equal to its length3, cubed. The SI unit of length, m, raised to the third power, is the derived SI unit of volume. A cube with an edge length of one meter has a volume of one cubic meter or m3.

• Chemistry uses smaller units like cm^3 (cc). The liter (L) is larger than a quart and equals a cubic decimeter (dm^3) in chemistry. The liter is our first non-SI metric unit. In a liter, 1000 milliliters (mL) equal 1 cm^3.

Common volumetric glassware:

Density

• Density is defined as the amount of mass in a unit volume of a substance:

• density=mass/volume

The terms density and weight are sometimes confused. A person who says that iron weighs more than air generally means that iron has a higher density than air—1 kg of air has the same mass as 1 kg of iron, but the iron occupies a smaller volume, thereby giving it a higher density. If we combine two liquids that do not mix, the less dense liquid will float on the denser liquid.

Units of Energy

• The SI unit for energy is the joule (pronounced “jool”), J, in honor of James Joule (1818−1889), a British scientist who investigated work and heat.

• Kilojoules (kJ) are used to discuss chemical reaction energies because joules are small amounts of energy. Hydrogen and oxygen react to generate 1 g of water, releasing 16 kJ of heat.

• A calorie (cal) was originally defined as the amount of energy required to raise the temperature of 1 g of water from 14.5 to 15.5 °C. It has since been defined in terms of a joule: 1 cal = 4.184 J (exactly).

• A related energy unit that is familiar to anyone who has read a food label is the nutritional Calorie (note the capital C), which is 1000 times larger than calorie with a lowercase c: 1 Cal = 1000 cal = 1 kcal.

Uncertainty in Measurement

• Two kinds of numbers are encountered in scientific work: exact numbers (those whose values are known exactly) and inexact numbers (those whose values have some uncertainty).

• Exact numbers can also result from counting objects. For example, we can count the exact number of marbles in a jar or the exact number of people in a classroom.

Precision and Accuracy

• The terms precision and accuracy are often used in discussing the uncertainties of measured values. Precision is a measure of how closely individual measurements agree with one another. Accuracy refers to how closely individual measurements agree with the correct, or “true,” value.

• The precision of the measurements is often expressed in terms of the standard deviation.

Significant Figures

• All digits of a measured quantity, including the uncertain one.

• To determine the number of significant figures in a reported measurement, read the number from left to right, counting the digits starting with the first digit that is not zero. In any measurement that is properly reported, all nonzero digits are significant.

• Zeros between nonzero digits are always significant—1005 kg (four significant figures); 7.03 cm (three significant figures).

• Zeros at the beginning of a number are never significant; they merely indicate the position of the decimal point—0.02 g (one significant figure); 0.0026 cm (two significant figures).

• Zeros at the end of a number are significant if the number contains a decimal point—0.0200 g (three significant figures); 3.0 cm (two significant figures).

Significant Figures in Calculations

• For addition and subtraction, the result has the same number of decimal places as the measurement with the fewest decimal places.

• For multiplication and division, the result contains the same number of significant figures as the measurement with the fewest significant figures. When the result contains more than the correct number of significant figures, it must be rounded off.

• In determining the final answer for a calculated quantity, exact numbers are assumed to have an infinite number of significant figures. Thus, when we say, “There are 12 inches in 1 foot,” the number 12 is exact, and we need not worry about the number of significant figures in it.

DIMENSIONAL ANALYSIS

• In dimensional analysis, units are multiplied together or divided into each other along with the numerical values. Equivalent units cancel each other. Using dimensional analysis helps ensure that solutions to problems yield the proper units. Moreover, it provides a systematic way of solving many numerical problems and of checking solutions for possible errors.

Conversion Factors

• The key to using dimensional analysis is the correct use of conversion factors to change one unit into another. A conversion factor is a fraction whose numerator and denominator are the same quantity expressed in different units.

Using Two or More Conversion Factors

• The first conversion factor is used to cancel meters and convert the length to centimeters. Thus, meters are written in the denominator and centimeters in the numerator. The second conversion factor is used to cancel centimeters and convert the length to inches, so it has centimeters in the denominator and inches, the desired unit, in the numerator.

Conversions Involving Volume

• The conversion factors previously noted convert from one unit of a given measure to another unit of the same measure, such as from length to length. We also have conversion factors that convert from one measure to a different one.