1. Unit 2: Biology (Chemistry Concepts) Objectives: I can … Identify common elements or compounds by their symbols or formulas. Write and read chemical equations and understand how/why reactions occur Differentiate between atomic mass & atomic number; between atoms, molecules, ions, & isotopes; between buffers, acids & bases; between protons, electrons, & neutrons; between carbohydrates, lipids, proteins, nucleic acids, & enzymes; between solute, solvent, solution, suspension, concentration, etc. Distinguish between types of bonds (Ex: covalent, ionic, hydrogen, etc) and their characteristics. Relate the unique properties of water to polarity, hydrogen bonds, gravity, and molecular mass Analyze the relationship between states of matter and kinetic energy. Compare & contrast the types of energy

2. Vocabulary: Matter, mass, atom, nucleus, neutron, proton, electron, element, atomic number, atomic mass, chemical symbol, isotopes, compounds, chemical formula, chemical equation, ions, ionic bonds, covalent bonds, molecules, structural formula, polarity, cohesion, adhesion, mixture, solution, isotonic, hypotonic, hypertonic, solute, solvent, concentration, suspension, colloid, hydronium ion, acid, hydroxide ion, base, alkaline, pH, energy (potential, kinetic), reactants, products, activation energy, exothermic, endothermic, monomer, polymer, carbohydrate, lipid, protein, enzyme, nucleic acid, substrate

3. matter - Occupies space and has mass mass - The amount of matter, or material, in an object (How is this related to weight?) atom - Basic unit of matter. Each type of atom has a unique number of subatomic particles (called protons, neutrons, and electrons) that determine its structure and behavior. Subatomic particles: 1) neutrons - have no charge, weigh approx. 1 amu, are in the nucleus (center) of the atom 2) protons - have a +1 (positive) charge, weigh 1 amu, and are also in the nucleus 3) electrons - have a -1 (negative) charge, weigh almost nothing, and move extremely rapidly at great distances from the nucleus. There is an equal number of protons and electrons in an atom so an atom has no overall charge.

5. Elements are a single type of atom. An element may be more than one atom of the same type. (Ex: elemental oxygen is naturally found as O 2 - two atoms of oxygen) Each element is listed in the periodic table and is given a chemical symbol, an atomic number, and an atomic mass number. Chemical symbol - a unique abbreviation (similar to initials) for each element. This is used in formulas and equations. Atomic number - the number of protons (positive charge) in an atom. Since atoms always have the same number of protons and electrons, this also tells us how many electrons are present. Atomic mass - the number of protons plus the number of neutrons (This is not usually a whole number since neutrons weigh slightly more than protons.)

6. 6 C Carbon 12.011 Section 2-1 An Element in the Periodic Table Go to Section:

8. Examples of commonly used chemical symbols: C = carbon Ca = calcium Na = sodium (natrium)K = potassium (kalium) H = hydrogenHe = helium Cl = chlorineMg = magnesium *Note: Only the first letter is capitalized for a chemical symbol. If we have 2 capitalized letters next to each other, we have 2 separate atoms and are looking at a chemical formula. Example: Co = cobalt (a single element) CO = carbon monoxide (a molecule with 2 types of atoms) To determine the number of atoms of each element in a chemical formula, we look at the subscripts. If no subscript, there is one atom of that element. Ex: CO = 1 carbon, 1 oxygen CO 2 = 1 carbon and 2 oxygen atoms

9. Chemical formulas show how many and what kinds of atoms are present in a molecule or compound. Ex: MgCl 2 shows there are 2 chlorine atoms for every magnesium atom. C 6 H 12 O 6 shows there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. CO shows there is 1 carbon and 1 oxygen (Co is cobalt) Molecules are atoms held together because they share electrons. This is a covalent bond. Ionic compounds are held together because 1 atom has a stronger pull for electrons than the other so it pulls the electron(s) over making it slightly negatively charged and the other atom slightly positively charged so they attract like magnets. This is an ionic bond.

10. Covalent bonds:

11. Ionic bonds:

12. Sodium atom (Na)Chlorine atom (Cl)Sodium ion (Na + )Chloride ion (Cl - ) Transfer of electron Protons +11 Electrons -11 Charge 0 Protons +17 Electrons -17 Charge 0 Protons +11 Electrons -10 Charge +1 Protons +17 Electrons -18 Charge -1 Section 2-1 Figure 2-3 Ionic Bonding Go to Section:

13. A structural formula shows which atoms bond where and by sharing or “stealing” how many electrons. Ex: Chemical formula: CO 2 Structural formula: O = C = O This shows carbon is between the two oxygen atoms and that 2 electrons are shared between each oxygen and the carbon. All atoms (except Helium and the very large atoms) want 8 electrons in their outer energy levels of electrons. This stabilizes them. This is the reason why chemical bonds form. The first energy level is stable with 2 electrons (He only has this energy level present). The next energy level needs 8 electrons but the first level must always be filled before going to the next level. So an atom with 10 electrons is stable (2 in first energy level and 8 in next.) An atom with 18 electrons is also stable (2, 8, 8).

14. When combining substances, we can use a chemical equation to show what atoms, etc. are present and how many of each. The chemicals we are adding together are called the “reactants.” The result of the combination is called the “product.” Ex: 2 Na + Cl 2  2 NaCl *Note: an arrow is used instead of an equal sign in chemical equations. But, like in math, the number of each “variable” (think of the elements added as “x’s” and “y’s”) must be equal on each side of the equation. The coefficient “2” in front of sodium (Na) tells us 2 separate atoms of sodium were part of the reactants. The subscript 2 on the Cl tells us a single molecule of chlorine (containing 2 atoms of Cl) was the other reactant. Note: a subscript only reflects back to the element immediately in front of it unless parentheses are used.

15. 2 Na + Cl 2  2 NaCl The product, 2 NaCl, tells us 2 molecules (EACH molecule containing one Na and one Cl ) were produced. Looking at other examples of chemical formulas using subscripts: MgCl 2 (1 magnesium, 2 chlorine atoms) Mg(HCO 3 ) 2 (1 Mg atom, 2 molecules of bicarbonate {HCO 3 } each containing 1 hydrogen, 1 carbon, and 3 oxygen atoms) Since the 2 is directly outside the parentheses, it distributes to everything within the parentheses. However, the 3 is next to the “O” and no parentheses is in front of it so it only applies to the oxygen. Total atoms: 1 Mg, 2 H, 2 C, 6 O How many of each type of atom do I have if I write 6 Mg(HCO 3 ) 2 ? 6 Mg, 12 H, 12 C, 36 Oxygen

16. Nonradioactive carbon-12Nonradioactive carbon-13Radioactive carbon-14 6 electrons 6 protons 6 neutrons 6 electrons 6 protons 8 neutrons 6 electrons 6 protons 7 neutrons Section 2-1 Figure 2-2 Isotopes of Carbon Go to Section:

17. Isotopes are atoms that have the same number of protons and electrons as the common atom of that type BUT have a different number of neutrons. These isotopes are often unstable and radioactive. Ex: Carbon 12 is stable with 6 neutrons (and 6 protons). The isotope, Carbon 14, has 8 neutrons and 6 protons. It breaks down the extra neutrons over time (radioactivity). Carbon dating measures the ratio of carbon 14 to carbon 12. Ions are atoms that have become charged (usually in a solution) by gaining or losing electrons. Ex: NaCl sometimes breaks into separate Na+ and Cl- ions in solution. Chlorine gains an electron (it already had 7 electrons in its outer energy level so it needed 1 to complete the level) and becomes negatively charged. Since sodium only had 1 electron in its outer energy level, it loses an electron to become positively charged.

18. Even water can ionize. H 2 O molecules “dissociate’ to become H 3 O+ cations (positively charged ions) and OH - anions (negatively charged ions). H 3 O+ is called a hydronium ion OH - is called a hydroxide ion When a solution has an equal number of hydronium (like pure water) and hydroxide ions, it is neutral. A solution with extra hydronium ions is acidic. A solution with fewer hydronium ions than hydroxides is basic, or alkaline. Ex 1: HCl + H 2 O  H 3 O + (aq) + Cl - (aq) (solution is acidic) Ex 2: NaOH + H 2 O  Na + (aq) + OH - (aq) (solution is basic)

19. Acids are proton donors. Acids tend to taste sour (Ex: lemon juice, vinegar) Many chemical formulas for acids are written starting with an “H” (the proton that is donated comes from hydrogen) Ex: HCl, H 2 SO 4, HCO 3 Bases are proton acceptors. Bases tend to taste bitter. Basic solutions are called alkaline. Strong bases are just as caustic (corrosive, burning) as strong acids. An acid reacted with a base can create a “salt” and water. Ex: NaOH + HCl  NaCl + H 2 O Acidity is measured on a pH scale. pH 7 = neutral, pH 0 = highly acidic pH 14 = highly alkaline Buffers help bring a solution closer to neutral.

22. Water Water is VERY unique. Water is polar. That is, the oxygen atom in water pulls harder on the electrons than the hydrogens do. So, the oxygen “edge” is more negative than the hydrogen portions. This makes water almost magnetic, so it likes to cling to surfaces. This is called adhesion. Because water is polar, ionic substances, like NaCl, dissolve easily in it. Nonpolar substances, like fat, don’t dissolve in water. Water forms drops and bubbles up on surfaces because it likes to cling to itself. This is called cohesion. Water molecules form hydrogen bonds BETWEEN MOLECULES, (Covalent and ionic bonds occur within a molecule or compound) causing water molecules to form groups of molecules instead of remaining as individuals. Why is this important? Because individual water molecules are lighter than air so water could not exist as a liquid if it didn’t group together! (Air is mostly N 2 gas) As temperature drops to freezing, water sets up a crystalline structure that spreads the molecules out a bit when forming ice. This makes water LESS DENSE as a solid! It is the only Earthly substance that is less dense as a solid.

23. Hydrogen bonds between water molecules (covalent bonds within water molecules).

24. Mixtures, Solutions and Suspensions A mixture - 2 or more elements or compounds physically mixed together but not chemically combined. Ex: salt mixed with sand When something like NaCl (salt) is mixed in water and the salt breaks into Na+ and Cl - ions, they equally disperse. This is a solution. The salt and water do not recombine to make a new substance such as NaOH and HCl so this is still a mixture. The taste of salt is still distinct letting us know it didn’t chemically react. The water is called a solvent. (Because it is polar, water can dissolve ionic and polar substances easily.) Solvents dissolve other things (Alcohol is nonpolar so it can dissolve nonpolar substances. In chemistry, like dissolves like) Concentration – amount of solute dissolved in solvent The salt, or anything being dissolved into something else, is the solute. Suspensions occur when molecules don’t dissolve by breaking into ions but are small enough to remain suspended within the liquid if it is stirred or moved often enough. If the mixing action stops for too long, the substance will settle out. Ex: Blood cells in blood plasma. “Shake well ” Colloids, or colloidal suspensions, don’t settle out even if not stirred. Ex: milk

25. Organic Chemistry Organic chemistry is also called the chemistry of carbon. Carbon has 4 electrons in its outer energy layer. This means it needs to either gain or lose 4 electrons to be happy. It pulls on its own electrons hard enough that they can’t be stripped away yet it isn’t strong enough to “steal” electrons from other atoms. So carbon shares (covalently bonds) easily with other atoms, sometimes sharing 1, 2, or 3 electrons with another atom. (In other words, forming single, double, or triple bonds.) Carbons can bond together to form chains, branched chains, or rings. Carbon can form large molecules, called macromolecules, that are essential to life. These include: Carbohydrates (simple and complex sugars) Lipids (fats, oils, waxes, steroids - think “cholesterol”) Proteins (includes enzymes, hormones, etc.) Nucleic acids (DNA, RNA)

26. MethaneAcetyleneButadieneBenzeneIsooctane Section 2-3 Figure 2-11 Carbon Compounds Go to Section:

27. Carbohydrates are made of carbon, hydrogen, and oxygen, usually in a ratio of about 1:2:1 Ex: C 6 H 12 O 6 (glucose) Carbohydrates with fewer carbons are called simple sugars. These include monosaccharides (the building blocks, or monomers, basic unit of sugars) and disaccharides. Most simple sugars are given names that end in “-ose.” Ex: glucose, fructose, lactose, dextrose, maltose, etc. Complex sugars (polymers) like starch and glycogen are called polysaccharides. (polymers are more than one monomer joined together) Simple sugars generally dissolve in water but complex sugars do not. --------------------------------- Lipids are also made of carbon, hydrogen, and oxygen but have far fewer oxygen atoms than carbohydrates. Fatty chains of carbon atoms with many double bonds between carbons are called polyunsaturated fats. Chains with all single bonds are called saturated fats. Polyunsaturated fats are usually liquid at room temperature (Ex: corn oil) Saturated fats like butter, animal fat, etc. are solids at room temperature. Every cell in our body has a membrane made of lipids!

28. Starch Glucose Section 2-3 Figure 2-13 A Starch Go to Section:

30. Proteins are made of C, H, O, and nitrogen that form amino acids, or protein monomers (building blocks) (amine = NH2, carboxylic acid = -COOH) Many different types of amino acids join via peptide bonds to form proteins. Some amino acids are polar and some are nonpolar. All proteins have a natural shape, or conformation. Some are folded, some form complex groups of subunits, etc. This natural shape is important to protein function. Changes in protein shape can damage or destroy its ability to perform. Heat, radiation, etc. can denature (unfold and destroy a protein’s function). Hormones, muscles, most bodily structures, enzymes, etc. are made of proteins. Enzymes are special proteins that help a chemical reaction take place, much like heat helps cake batter become an actual cake. The heat is not an actual ingredient in the cake, it doesn’t chemically react with the ingredients, yet without it, no cake. This is how enzymes behave in chemical reactions. They lower the “activation energy” to get a reaction going. Enzymes that speed up reactions are called catalysts. The reactants in an enzyme “catalyzed” reaction are called substrates.

31. Reaction pathway without enzyme Activation energy without enzyme Activation energy with enzyme Reaction pathway with enzyme Reactants Products Section 2-4 Effect of Enzymes Go to Section:

32. General structureAlanineSerine Section 2-3 Figure 2-16 Amino Acids Go to Section: Amino groupCarboxyl group

33. Nucleic acids contain C, H, O, N, and phosphorus monomers known as nucleotides. Each nucleotide contains a 5 carbon sugar ring (ribose or deoxyribose), a phosphate (PO4-) group, and a nitrogenous base (adenine, guanine, cytosine, thymine, uracil). DNA (deoxyribonucleic acid) is found in the nucleus of most cells and controls their functions as well as determining our inherited traits. RNA (ribonucleic acid) helps carry out the DNA instructions to make proteins, etc. Interestingly, ATP (our cells’ energy source) is very similar to RNA and DNA. It contains a 5 carbon sugar, the nitrogenous base, adenine, and 3 phosphate groups. Suggestion: Make a concept map for macromolecules.

35. Carbon Compounds Macromolecules include Examples which contain Examples which contain Carbon, hydrogen, oxygen Carbon, hydrogen, oxygen

36. Energy plays an important role in everything. Potential energy is the “ability” to do work (movement of mass). It is stored energy or energy due to position. (Ex: fat is stored energy, a battery not in use has potential energy, a rock at the edge of a cliff has potential energy, chemical bonds have potential energy) Kinetic energy is energy of motion. (Current running through a wire is kinetic energy, movement is kinetic energy.) It is important to remember that kinetic energy is what makes a substance a gas, liquid, or solid. The faster the molecules move in a substance, the more kinetic energy it has. Solids have molecules that are barely moving, just vibrating in place. (Low kinetic energy) Liquids have molecules bouncing off each other and sliding past one another. (Medium kinetic energy) Gases have molecules rapidly moving, slamming against each other and bouncing far distances apart. (High kinetic energy) Heating a substance can increase its kinetic energy.

37. Energy can be converted from one form to another. Ex: Mechanical to electrical (generators, alternators); chemical to heat & light (burning a match), solar to chemical (photosynthesis) Molecules can hold energy (potential energy) in their bonds. Breaking or forming bonds during a chemical reaction can consume (absorb) or release energy, often in the form of heat. Activation energy is the energy needed to get a reaction started. Reactions that give off more heat energy than they absorb are called exothermic reactions. Reactions that use more heat energy than they produce are called endothermic reactions.