Prep for chemistry entrance exams the right way. Review summary concepts, watch animations, and take unlimited practice exams. This animation analyses and interprets the sp3 hybridization in carbon, taking ethane as the reference example.
1. Matter takes up space and has mass.
2. All living and nonliving matter is composed of 92 naturally-occurring basic elements.
3. Elements cannot be broken down to substances with different chemical or physical properties.
4. Six elements (C, H, N, O, P, S) make up 98% of living things.
1. Chemical and physical properties of atoms (e.g., mass) depend on the subatomic particles.
a. Different atoms contain specific numbers of protons, neutrons, and electrons.
b. Protons and neutrons are in nucleus of atoms; electrons move around nucleus.
c. Protons are positively charged particles; neutrons have no charge; both have about 1 atomic mass
unit of weight.
d. Electrons are negatively charged particles.
2. The atomic mass of an atom is about equal to the sum of its protons and neutrons.
3. All atoms of an element have the same number of protons, the atom’s atomic number.
The Periodic Table
1. The periodic table shows how various characteristics of atoms recur.
2. The table is arranged in order of atomic number, with periods in horizontal rows and groups in vertical columns.
1. Isotopes are atoms with the same number of protons but differ in number of neutrons; e.g., a carbon
atom has six protons but may have more or less than usual six neutrons.
2. A carbon with eight rather than six neutrons is unstable; it releases rays and subatomic particles and is
a radioactive isotope.
3. Low levels of radiation such as radioactive iodine or glucose allow researchers to trace the location and
activity of the atom in living tissues; therefore these isotopes are called “tracers.”
4. High levels of radiation can cause cancerous tissues and destroy cells; careful use of radiation in turn
can sterilize products and kill cancer cells.
Electrons and Energy
1. Electrons occupy an orbital at some level near or distant from the nucleus of the atom.
2. An orbital is a volume off space where an electron is most likely to be found; an orbital contains no more than two electrons.
3. The more distant the orbital, the more energy it takes to stay in the orbital.
4. When atoms absorb energy during photosynthesis, electrons are boosted to higher energy levels.
5. The innermost shell of an atom is complete with two electrons; all other shells are complete with eight electrons.
1. When two or more different elements react or bond together, they form a compound (e.g., H2O).
2. A molecule is the smallest part of a compound that has the properties of the compound.
3. Electrons possess energy and bonds that exist between atoms in molecules contain energy.
1. Ionic bonds form when electrons are transferred from one atom to another.
2. Losing or gaining electrons, atoms participating in ionic reactions fill outer shells, and are more stable.
3. Example: sodium with one less electron has positive charge; chlorine has extra electron that has
negative charge. Such charged particles are called ions.
4. Attraction of oppositely charged ions holds the two atoms together in an ionic bond.
1. Covalent bonds result when two atoms share electrons so each atom has octet of electrons in the outer
2. Hydrogen can give up an electron to become a hydrogen ion (H+
) or share an electron with another
atom to complete its outer shell of two electrons.
3. Structural formulas represent shared atom as a line between two atoms; e.g., single covalent bond
(H–H), double covalent bond (O=O), and triple covalent bond (N = N).
4. Three dimensional shape of molecules is not represented by structural formulas but shape is critical in
understanding the biological action of molecules: action of insulin, HIV receptors, etc.
Nonpolar and Polar Covalent Bonds
1. In nonpolar covalent bonds, sharing of electrons is equal.
2. With polar covalent bonds, the sharing of electrons is unequal.
a. In water molecule (H2O), sharing of electrons by oxygen and hydrogen is not equal; the oxygen
atom with more protons dominates the H2O association.
b. Attraction of an atom for electrons in a covalent bond is called electronegativity; oxygen atom is
more electronegative than hydrogen atom.
c. Oxygen in water molecule, more attracted to electron pair, assumes small negative charge.
1. A hydrogen bond is weak attractive force between slightly positive hydrogen atom of one molecule
and slightly negative atom in another or the same molecule.
2. Many hydrogen bonds taken together are relatively strong.
3. Hydrogen bonds between complex molecules of cells help maintain structure and function.
Chemistry of Water
1. All living things are 70–90% water.
2. Because water is a polar molecule, water molecules are hydrogen bonded to each other.
3. With hydrogen bonding, water is liquid between 0º C and 100 C which is critical for life.
4. The temperature of liquid water rises and falls more slowly than that of most other liquids..
a. Calorie is amount of heat energy required to raise temperature of one gram of water 1º C.
b. Because water holds more heat, its temperature falls more slowly than other liquids; this protects
organisms from rapid temperature changes and helps them maintain normal temperatures.
Heat of Vaporization
a. Hydrogen bonds between water molecules require a large amount of heat to break.
b. This property moderates earth’s surface temperature; permits living systems to exist here.
c. When animals sweat, evaporation of the sweat takes away body heat, thus cooling the animal.
a. Cohesion allows water to flow freely without molecules separating, due to hydrogen bonding.
b. Adhesion is ability to adhere to polar surfaces; water molecules have positive, negative poles.
c. Water rises up tree from roots to leaves through small tubes.
d. Adhesion of water to walls of vessels prevents water column from breaking apart.
e. Cohesion allows evaporation from leaves to pull water column from roots.
Acids and Bases
1. Covalently bonded water molecules ionize; the atoms dissociate into ions.
2. When water ionizes or dissociates, it releases a small (107 moles/liter) but equal number of H+ and OH ions; thus, its pH is neutral.
3. Water dissociates into hydrogen and hydroxide ions: H – O –H → H+ + OH-
4. Acid molecules dissociate in water, releasing hydrogen ions (H+) ions: HCl → H+ + Cl-
5. Bases are molecules that take up hydrogen ions or release hydroxide ions. NaOH → Na+ + OH-
6. The pH scale indicates acidity and basicity (alkalinity) of a solution.
a. Measure of free hydrogen ions as a negative logarithm of the H+ concentration (-log [H+]).
b. pH values range from 0 (100 moles/liter; most acidic) to 14 (1014 moles/liter; most basic).
1) One mole of water has 107 moles/liter of hydrogen ions; therefore, has neutral pH of 7.
2) Acid is a substance with pH less than 7; base is a substance with pH greater than 7.
3) As logarithmic scale, each lower unit has 10 times the amount of hydrogen ions as next higher pH unit; as move up pH scale, each unit has 10 times the basicity of previous unit.
7. Buffers keep pH steady and within normal limits in living organisms.
a. Buffers stabilize pH of a solution by taking up excess hydrogen (H+) or hydroxide (OH-) ions.
b. Carbonic acid helps keep blood pH within normal limits: H2CO3 → H+ + HCO3-.
1. Carbon has four electrons in outer shell; bonds with up to four other atoms.
2. Addition of an –OH (hydroxyl group) turns a carbon skeleton into an alcohol.
3. The ethanol alcohol is hydrophilic, it dissolves in water, because the –OH functional group is polar.
4. A hydrocarbon is hydrophobic except when it has an attached ionized functional group such as
carboxyl (acid) (— COOH); then the molecule is hydrophilic.
5. Carboxyl groups ionize in solution and release hydrogen ions, being both polar and acidic.
6. Isomers are molecules with identical molecular formulas but differ in arrangement of their atoms.
1. Polymers are the large macromolecules composed of three to millions of monomer subunits.
2. Polymers build by different bonding of different monomers; mechanism of joining and breaking these
bonds is dehydration synthesis and hydrolysis.
3. Cellular enzymes carry out condensation synthesis and hydrolysis of polymers.
4. During dehydration synthesis, a water is removed and a bond is made (synthesis).
5. Hydrolysis reactions break down polymers in reverse of dehydration; a hydroxyl (— OH) group from
water attaches to one monomer and hydrogen (— H) attaches to the other.
1. Monosaccharides are simple sugars with a carbon backbone of three to seven carbon atoms.
2. Glucose and fructose isomers have same formula (C6H12O6) but differ in structure.
3. Ribose and deoxyribose are five-carbon sugars (pentoses); the backbones of DNA.
4. Disaccharides contain two monosaccharides joined by dehydration synthesis.
5. Polysaccharides are polymers of monosaccharides.
6. Starch is straight chain of glucose molecules with few side branches.
1. Lipids contain two molecular units: glycerol and fatty acids.
2. Glycerol is a water-soluble compound with three hydroxyl groups.
3. Triglycerides are glycerol joined to three fatty acids by dehydration synthesis.
4. A fatty acid is a long hydrocarbon chain with a carboxyl (acid) group at one end.
a. Most fatty acids in cells contain 16 to 18 carbon atoms per molecule.
b. Saturated fatty acids have no double bonds between their carbon atoms.
c. Unsaturated fatty acids have double bonds in the carbon chain where there are less than two hydrogens per carbon.
1. Phospholipids are like neutral fats except the third fatty acid is replaced by phosphate group or a
group with both phosphate and nitrogen.
2. The phosphate group bonds to another organic group (R) and is the polar head; hydrocarbon chains
become nonpolar tails.
3. Phospholipids arrange themselves in a double layer in water, so the polar heads face outward toward
water molecules and nonpolar tails face toward each other away from water molecules.
Free Practice Tests
Each of the following multiple-choice chemistry tests has 10 questions, reviewing your knowledge of basic chemistry and principles. No sign up required, just straight to the test.
The Periodic Table
Properties of alkali, alkaline earth and transition metals. Halogens and noble gases. Atomic weights and elements.
Reactants and products in reversible and irreversible chemical reactions, just like your own chemistry tutor in the lab.
Oxidation and Reduction
Introducing oxidation states, oxidation, and reduction. Some tips for remembering oxidation and reduction.
Net Ionic Equations
Net Ionic Equations; Representations of Reactions; Physical and Chemical Changes. Determine a balanced chemical equation for a given chemical phenomenon.
Try clicking on one of the links below. This is an excellent reference list to study from, as it returns visual details and the mechanism of both the specified chemical reaction, as well as related reactions.
Chemistry Compound Database - searchable by partial name, melting point, boiling point, index of refraction,
molecular formula, UV absorption, mass spectral peaks, and chemical types.
Chemical engineers apply the principles of chemistry, biology, physics, and math to solve problems that involve the production or use of chemicals, fuel, drugs, food, and many other products. They design processes and equipment for manufacturing, test production methods, and oversee byproducts treatment. Process engineers are specialists working in a particular process, such as oxidation or polymerization (making plastics and resins).
Chemists may specialize in a particular field, such as nanomaterials or biological engineering. Still others specialize in developing specific products. In this capacity, chemical engineers may be exposed to health or safety hazards when handling certain chemicals and plant equipment, but such exposure can be minimized if proper health and safety procedures are followed.
What do chemical engineers do on a daily basis?
- Conduct research to develop new and improved manufacturing processes.
- Establish safety procedures for those working with dangerous chemicals.
- Develop processes for separating components of liquids and gases, or for generating electrical currents, by using controlled chemical processes.
- Design and plan the layout of equipment.
- Conduct tests and monitor the performance of processes throughout production.
- Troubleshoot problems with manufacturing processes.
- Evaluate equipment and processes to ensure compliance with safety and environmental regulations.
- Estimate production costs for management.
Chemistry Job Listings
Employment of chemical engineers is projected to grow 8 percent from 2021 to 2025. Many chemical engineers work in industries whose products are sought by many manufacturing firms. For instance, they work for firms that manufacture plastic resins, which are used to increase fuel efficiency in automobiles. Increased availability of domestically produced natural gas should increase manufacturing potential in the industries employing these engineers. In addition, chemical engineering will continue to migrate into dynamic fields, such as nanotechnology, alternative energies, and biotechnology, and thereby help to sustain demand for engineering services in many manufacturing industries.
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