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Atomic Structure • What is an atom? • An atom is the smallest unit of matter. It contains protons (+ charge), neutrons (0 charge), and electrons (- charge). The protons and neutrons are located in the nucleus of the atom. Electrons orbit around the nucleus in distinct energy levels. • The nucleus of the atom is small, dense, positively charged, and heavy. Ernest Rutherford determined this during his gold foil experiment. Rutherford Experiment Rutherford Experiment Atomic Structure • The nucleus is very small compared to the rest of the atom. Its size is comparable to a marble resting at the 50 yard line of Qualcomm Stadium. Therefore, the atom is mostly made up of empty space. • Even though like charges repel, the protons in the nucleus do not repel because they interact with and are spaced out by the neutrons. Protons, Neutrons, and Electrons • How do we find the numbers of each subatomic particle? • The number of protons=atomic number (the whole number on the periodic table). Do not confuse atomic number with atomic mass, they are two different things. • #protons=#electrons in an atom with no charge. # neutrons • Mass number= #protons + #neutrons. This can be rearranged to determine neutrons. How can we do it? • The mass number cannot be found on the periodic table. The atomic mass is the average of an element’s mass numbers, but not the mass number we use for problems. • How many protons, neutrons, and electrons are in phosphorous 31? (mass number=31) Atom Example QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Problems • How many protons, neutrons and electrons are in Oxygen 16 (mass number =16)? • How many protons, neutrons, and electrons are in Iron 56 (mass number=56)? What are isotopes? • Isotopes are atoms of the same element that have different numbers of neutrons. • Some isotopes are stable and some are not. The unstable isotopes decay and we call them radioactive. Unstable isotopes decay at a known rate and allow us to date artifacts and rocks using that known rate. • Consider Carbon 12, Carbon 13, and Carbon 14. Are these isotopes? How many protons, neutrons, and electrons are found in each? • I will show you isotope notation on the board also. Isotope Examples QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are periodic families? • The rows on the periodic table are called periods. Each period represents an energy level. The columns are called groups. • The families on the periodic table occur in columns (up and down) and contain elements that have the same physical and chemical properties. • Elements that are to the right of the staircase are nonmetals. Elements that are to the left of the staircase are metals. • Elements in column one are called “alkali metals.” • Elements in column two are called “alkaline earth metals.” • Elements in columns 3-12 are called “transition metals.” • Elements in column 17 are called “halogens” • Elements in column 18 are called “noble gases” or “inert gases.” • Silicon and Germaniuim are “metalloids” meaning that they have properties of both metals and non-metals. • What families do these elements belong in? Questions • Name an element that is a member of: The halogens The inert gases The transition metals The alkali metals • Mg is a member of what family? What are the periodic trends? • When we look at periodic trends, across means we go from left to right and down a group means that we go from top to bottom. These trends are not readily obvious from the periodic table. • Atomic size/ionic size- This is the size of the atom or ion. This decreases across a row and increases when you go down a group. • Electronegativity-this is the tendency of an element to attract electrons. This increases across a row and decreases down a group. • Ionization energy-this is the amount of energy required for an atom to eject one of its electrons. This increases across a row and decreases down a group. • Between Al, P, and Cl, which element has the higher electronegativity? Questions • Between N, P, and As, which element has the highest atomic size? • Between C, N, and O, which element has the highest ionization energy? What is radioactive decay? • Radioactive decay occurs when an unstable isotope decays into smaller isotopes. • There are three types of radioactive decay. Alpha ()-this is when a helium nucleus is ejected from an isotope. This is the weakest form of radiation. Paper and clothing can stop it. • Beta()-this occurs when a neutron emits an electron. This neutron becomes a proton. This is middle strength radiation and wood or cardboard can block it. Penetrating Power of Radiation QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Types of radiation • Gamma()- this is high energy radiation that a radioactive isotope can emit when it decays. It causes no change in the nucleus. However, it is the strongest form of radiation. Lead and concrete can block it. • For nuclear reactions, please look on the board. What is fission or fusion? • Fission is when a radioactive isotope splits apart into smaller components. A lot of energy and gamma radiation is released. • Fusion is when two hydrogen nuclei fuse together to create a helium nucleus. A lot of energy and gamma radiation is released (more than fission). • Fusion of higher elements occurred in dying stars. There was enough heat and pressure, within those dying stars, for the helium to fuse with other helium atoms and form larger elements. The destruction of that star then diffused those elements into the universe. • There was also enough energy for those elements, once they had formed, to fuse with other nuclei and form larger elements in dying stars. • Thus, all life on this planet is the product of dying stars. Examples of Fusion and Fission • Fission is first picture. • Fusion is in second picture. Quic kT i me™ and a T IFF (Unc ompres s ed) dec ompres s or are needed t o s ee thi s pi c ture. Quic kT i me™ and a T IFF (Unc ompres s ed) dec ompres s or are needed t o s ee thi s pi c ture. What are orbitals? • Orbitals are clouds where electrons reside in an atom. • Each orbital can hold a maximum of two electrons (one with up spin and one with down spin). • The types of orbitals are: s (1 orbital, energy level=row # on periodic table) p (3 orbitals, energy level=row # on periodic table) d (5 orbitals, energy level=row # on periodic table -1 ) f (7 orbitals, energy level=row # on periodic table -2 ) Orbital example Question? • How many electrons can each orbital hold? • Where are these orbital blocks located on the periodic table? What are electron configurations? • Electron configurations show how electrons are arranged around an atom. • Three rules govern electron configurations and they are: • Aufbau principle-electrons enter orbitals of lowest energy first. For all electron configurations, start filling electrons in the 1s orbital, which is the lowest energy level. • Pauli exclusion principle-An atomic orbital may describe at most two electrons (no more than two electrons per orbital) • Hund’s rule-When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with spins parallel. Second electrons then add to each orbital so that their spins are paired with those of the first electrons. How do we write electron configurations? • Electron configurations contain the energy level, orbital, and number of electrons in that orbital. For example, 1s1, the first number means energy level one, s is the orbital, and the superscript number means that there is one electron in this orbital. • Start at hydrogen and fill electrons from left to right starting from there. Count how many electrons are in these orbitals as you move from left to right. Also, make sure you change to the correct orbital once you move to another orbital’s block on the periodic table. • You can move to another row on the periodic table after you have moved to all the way over to the right on the previous row. Remember that energy levels change when you move to another row on the periodic table. Do this until you have all electrons accounted for in the element. Write electron configurations • Write the electron configuration for hydrogen. • Hydrogen has one electron. Begin at row one, s block, and count over one. That is hydrogen. • From that, we are in energy level one, the s orbital, and there is only one electron in it. So we write, 1s1 for the electron configuration. • What about for helium? Write its electron configuration. Orbital Blocks More electron configurations • Write the electron configuration for Li. • It should be: • 1s2 2s1 and this shows three total electrons, which lithium has. • Write the electron configuration for N • It should be: • 1s2 2s2 2p3 Write the electron configuration for Fe How do we write short-hand electron configurations? • Start short-hand electron configurations with the nearest noble gas that has fewer electrons than the element you are looking at. • Write that noble gas and its number of electrons once you begin your electron configuration. Then, write your electron configuration from that noble gas. • For example, for Mg, you would begin its short-hand electron configuration with [Ne]10 • That would take care of the first 10 electrons in that configuration (through row 2). You would then move to row three and complete the electron configuration. • It would look like: [Ne]10 3s2 Write the short-hand electron configuration for: • • • • Na Fe Br As What is another way to write electron configurations? • This method of writing electron configurations shows the orbital (1s for example) and how many electrons are in it. Arrows represent the electrons (one arrow points up meaning that the electron has up spin, and another arrow points down meaning that the other electron has down spin) and are placed in circles that represent the orbitals. • When writing electron configurations in this way, follow Hund’s rule meaning that you must fill each orbital, in a particular subshell, with one electron that has the same spin (usually an up arrow). After filling these orbitals, then you would place the second electron in each orbital and it would have down spin (opposite of the first electron). • For example, for the p orbital, you would place one electron (up arrow) in each of the three orbitals before filling each orbital with the second electron. The same holds true for the d orbitals. Example • Look on the whiteboard for examples of this style of electron configuration. • Write electron configurations for the following: Ne O Na How do we use the short-hand method? • For this type of electron configuration, you can use the short-hand method. You find the nearest inert gas (to the element) that has fewer electrons and place it in brackets at the beginning of the configuration. Then you write the configuration from there showing the orbitals and how many electrons are in them. • Write short-hand electron configurations for: Fe Br Sr What are valence electrons? • Valence electrons are electrons in the outermost energy level of an atom. They are the electrons that are involved in bonding of atoms. • Core electrons are electrons that are in the inner energy levels of an atom. They are not involved in bonding. • s and p electrons (if they are in the same highest energy level) automatically count as valence electrons. • d electrons count as valence electrons if there are 5 or fewer of them in the outermost energy level. If there are 6 or more d electrons then they do not count as valence electrons. How do we determine the number of valence electrons? • We must first find the highest energy level in an atom (look at the highest energy number in its electron configuration, this is easy to do if you use short-hand). Count the number of electrons in that highest energy level (s and/or p electrons if they are in the same highest energy level and d electrons if they meet the criteria from the last slide). • How many valence electrons does Mg have? • Electron configuration: [Ne]10 3s2 • Highest energy level is energy level 3 and there are only two electrons there, which means that Mg has 2 valence electrons. Find valence electrons for the following elements: • • • • • • Be O F Br As Fe What are the types of bonding? • Bonding is when two or more atoms come together. Valence electrons are involved in bonding. • Ionic bonding is the bonding between metals and non-metals. These compounds are solid at room temperature, most of them dissolve in water and they conduct electricity when they dissolve. • Metallic bonding is the bonding between metals. These compounds do not dissolve in water but they do conduct electricity at room temperature. Ionic Substance • Covalent bonding-this is the bonding between non-metals. Some of these substances dissolve in water. Some do not, but they never conduct electricity. • Covalent Network solids-These are hard solids made up of carbon or silicon atoms in a network. An example of a covalent network solid is … Covalent Substance-Water Covalent Network Solid • What types of bonding do the following substances exhibit? NaOH C(graphite) Cr CO How do ions form? • Ions form by losing or gaining electrons in order to get a noble gas electron configuration. • Metals lose electrons and get a positive charge. #electrons lost= +charge of metal • Non-metals gain electrons and get a negative charge. #electrons gained= charge of non-metal. Ionic Bonding • Ionic Bonding occurs between metals and non-metals when electrons are transferred. • In order to form ionic compounds, we must first form ions. • Ions are elements that have a charge because they lose or gain valence electrons. • Cations have a positive charge while anions have a negative charge. Ion formation QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Ions continued • For metals, metals lose valence electrons until they get a noble gas configuration and get a positive charge (#electrons lost=positive charge of element). • Non-metals gain valence electrons until they get a noble gas configuration and get a negative charge (#electrons gained=negative charge of element). • What is the charge of chlorine when it becomes an ion? What about magnesium, sodium, aluminum? How do we form ionic compounds? • We form ionic compounds by putting a metal ion together with a non-metal ion. We cannot change their charges but we can change the numbers of each. • We need to make ionic compounds that are electrically neutral. In order to do this, we need to add more of each ion until this occurs. • A quick way to do this, if their charges do not cancel, is to cross the coefficients (not the signs) of the charges. Make sure that you have the lowest whole number ratio between charges. Making ionic compounds • Make ionic compounds between: Na and Cl Al and F Mg and N What are polyatomic ions? • Polyatomic ions are ions that have multiple elements in them and they have an overall charge. Look on your half sheet for those charges. For polyatomic ions, place them in parentheses before you cross the charges. • Make ionic compounds with the following: Na and SO4 Mg and NO3 K and PO4 Network of NaCl ions • This shows how salt crystals are put together. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. How do we name ionic compounds? • Ionic compounds consist of two names that are based on the elements that they contain. The first name is always the metal that is in the compound (the element or polyatomic ion that you write first in the compound). The second name is for the non-metal in the compound. • If the non-metal is a singular ion, then you use the name of the element but you take off the last syllable and add -ide. For example, chlorine becomes chloride when it is in an ionic compound or oxygen becomes oxide. • If the non-metal is a polyatomic ion then you use the name of the polyatomic ion from your half sheet. • For example, CaCl2 has the name of Calcium Chloride. Questions • Name the following ionic compounds: NaCl Cu(NO3)2 AlBr3 FeO MgSO4 ZnS • Convert the following back into ionic compounds: Sodium Cyanide Aluminum Chloride Magnesium Oxide Zinc Nitrate How do we name covalent compounds? • For covalent compounds, you name both elements that are involved in the compound. You write the name of the first element and leave it unchanged. You then write the name of the second element, after the first element, take off its last syllable and add -ide. • You also must use prefixes with covalent compounds. These prefixes denote how many atoms of each element are in the compounds. The prefixes are: mono (1, use on second element only), di (2), tri(3), tetra (4), penta (5), hexa (6), hepta (7), octa (8), nona (9), and deca (10). • For example, CO2 has the name of carbon dioxide. • Name these covalent compounds: CCl4 NH3 SO2 N2O4 • Write covalent compounds from these names: Carbon Monoxide Dihydrogen Monoxide Dinitrogen Pentaoxide Covalent Bonding • Covalent bonding is between non-metals where electrons are shared. Large biological molecules have this type of bonding. • In order to demonstrate covalent bonding, we first must make Lewis dot symbols. • You make these by taking an element’s valence electrons and placing them around the element symbol as dots. You place one around each of the four corners and then add more as needed. Each side can only have two electrons. Covalent Bonding • Unpaired electrons (1 dot) represent available spots where covalent bonding can occur. • Paired up electrons represent lone pairs, where bonding cannot occur. • Make Lewis dot symbols for: Ne S F N C Lewis structures • Lewis structures are a 2-D representation of covalent molecules. • In order to make them, first split the molecule into its component elements. • Put Lewis dot symbols around each element. • Pair up unpaired electrons, on different atoms, to form covalent bonds (1 bond=2 shared electrons). • Put it all together. Do these examples: • HF • CH4 • PBr3 QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What is the octet rule? • The octet rule says that atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell. This means that all atoms, in a Lewis structure, must have eight valence electrons around them (they can be either bonded or lone pair electrons). • Hydrogen and helium are exceptions to the octet rule. There is one more element that is an exception to the octet rule, which one is it? • Some elements (like sulfur) can have more than eight electrons around them, which means that they also break the octet rule. Practice • Complete the following Lewis structures: C2H5OH CCl4 BH3 Multiple Bonds • In order to form multiple bonds, you must first make single bonds between the atoms. If there are still unpaired electrons then you can start forming double or triple bonds. • Double bond=4 electrons between the atoms. • Triple bond=6 electrons between the atoms. • Make the following molecules: O2 COHF HCN What do we need to know about biochem? • Carbon makes up the backbone of most biological molecules. Carbon’s ability to form four bonds (single, double, or even triple bonds) is responsible for this. • Large molecules (polymers), such as starch, proteins, and nucleic acids are formed by repetitive combinations of simple subunits. • Amino acids are the subunits and hence building blocks of proteins. More Lewis Structures • Draw Lewis Structures for the following: N2 COH2 CO2 What is polarity? • Polarity is the result of unequal sharing of electrons. A polar covalent bond results when there is a difference in electronegativity between the elements involved in that bond. • Due to this electronegativity difference, the bonding electrons are shared unequally. They usually spend more time around the more electronegative atom. This creates a partial positive and partial negative end of the atom. • A dipole is created from this and it goes from partial negative to partial positive end. How do we know if a bond is polar? • We must compare the difference in electronegativities, between the two elements, in order to determine whether or not the bond is polar. The electronegativity for each element is located on pg. 366. • In order to do this, we must subtract the electronegativity values for each element. We can classify the bond, based on that difference, using the table on pg. 418. We can classify the bond as non-polar, moderately polar, very polar, and ionic. • Classify the following bonds as non-polar, polar, or ionic: • H-F K and Cl Br-Br How do we know if a molecule is polar? • If a molecule is polar, there must be a net dipole. • If a molecule is diatomic, with the same element, then it is non-polar. If there are different elements then it is polar. • If a molecule contains only C-C or C-H bonds then the molecule is non-polar. • If a molecule contains a lone pair around the central atom then that molecule is polar. • If the central atom is carbon, and it is bonded to the same element all around, then the molecule is non-polar. • If the central atom is carbon and it is bonded to two or more different elements all around then the molecule is polar. • Polar or non-polar? CH4 NH3 COHF O2 Example of dipole on a polar molecule. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are intermolecular forces? • Intermolecular forces are forces between molecules. Substances with strong intermolecular forces will have a high boiling point and a low freezing point. • The types: Van Der Waal’s or London Dispersion forces-these are the forces between non-polar molecules. A temporary dipole forms between the molecules. Due to their lack of true dipoles, nonpolar substances cannot interact with polar substances, which is why they cannot dissolve in polar substances such as water. These forces are the weakest. • Dipole-Dipole-These are the forces between polar molecules. Their dipoles interact. The partial positive ends interact with the partial negative ends. These forces are middle strength. • Hydrogen Bonding-These are between polar substances that have a strong dipole, which forms when hydrogen is bonded to N, O, or F. These are the strongest behind ionic forces. • Ionic forces-forces between ionic compounds. These are very strong due to their charges. Example of Hydrogen Bonding QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. • Using intermolecular forces, which has the highest boiling point between H2O, H2S, CH4, and HCl? • How do ammonia (NH3) molecules interact? Can they dissolve in octane (C8H18)? What is a mole? • Moles are a counting unit that we use in chemistry. They allow us to convert between quantities. • 1 mole=22.4 L at STP • STP is standard temperature and pressure, which is 0 C and 1 atmosphere of pressure. • We do mole conversions using conversion factors based on the equivalences. The units that you want go on top and the units that you want to cancel go on the bottom. • Use your mole road map. Not the chemistry mole Problems • How many moles are in 44.8 L of CO2 at STP? • How many liters are in 2 moles of O2 gas at STP? What is Avogadro’s number? • Amadeo Avogadro, an Italian scientist, hypothesized that equal volumes of gas, at equal temperatures and pressures, contain the same number of particles. Scientists eventually proved him right. • Avogadro’s number is 6.02 x 1023 particles. These particles could be atoms (singular elements), molecules (covalent compounds), or formula units (ionic compounds). • 1 mole= 6.02 x 1023 particles (atoms, formula units, molecules) Avogadro’s hypothesis Problems-use mole road map • How many moles are in 5 x 1020 molecules of CO2? • How many atoms of Cu are in 2 mol of Cu? What is the gram formula mass? • The gram formula mass (gfm) of a substance is defined as the sum of all the atomic masses of the elements in a compound multiplied by the number of each respective element. • To find the gfm, find what elements and how many of each element you have in a compound. Multiply the number of each element by its atomic mass and add them up. • For example, in CaCl2, you have 1 Ca and 2 Cl. That corresponds to the following gfm: 40.08 (atomic mass of 1 Ca) + 2 x 35.5 (atomic mass of Cl x 2)=111.08 g Find the gfm for the following substances: • • • • NaCl Cu(NO3)2 MgSO4 CO2 What about moles and mass? • 1 mole = gfm of that element or compound. • This is the most important mole relationship and we will use it all year long. • How many moles are in 50 g of NaCl? • How many grams are in 2 moles of Zn(OH)2? • How many moles are in 20 g of MnO2? Mole Review • Remember that: 1 mole= 6.02 x 1023 particles 1 mole = 22.4 L of a gas at STP 1 mole= gfm of a substance Problems • How many moles are in 1.204 x 1024 formula units of NaCl? • How many liters are in 3 moles of O2 at STP? • How many moles are in 100 g of Pb(NO3)2? Mole Review Transparency • Convert the following to moles: 20 g NaCl 3 x 1024 atoms He 40 g AlPO4 10 g Ba(OH)2 44.6 L N2 4 x 1012 atoms Ca Convert the following from moles 2 mol CaSO4 to grams 3 mol He to atoms 5.6 mol N2 to L 2.3 mol Al(NO3)3 to grams 6 mol Zn to atoms 5 mol NaOH to grams What are two-step mole conversions? • Two-step mole conversions allow us to convert between two different units on the mole road map. They involve two conversions. The first conversion takes you from your initial quantity to moles. The second conversion takes you from moles back out to your wanted quantity. • In these two-step conversions, moles are the intermediate. • This is a very powerful tool. For example, it can tell us the mass of one molecule of carbon dioxide. Problems • How many grams are in 11.2 L of Ne gas at STP? • How many grams are in 1 molecule of Carbon Dioxide (CO2)? What is percent composition? • Percent composition is the percent by mass of each element in a compound. The percent mass of an element in a compound is the number of grams of the element (that make up the gfm of the compound) divided by the total gfm of the compound, multiplied by 100%. • %mass of element E= (grams of element in one mole of the compound)/gfm of compound) x 100 • That formula applies for all elements in the compound. Problems • Calculate the percent composition of each element in C3H8. • Calculate the percent composition of each element in NaNO3. What is the empirical formula? • The empirical formula of a compound gives the lowest whole number ratio of the atoms of the elements in a compound. • The empirical formula is valuable because it tells us the kinds and relative count (atoms or moles of each element) in molecules or formula units of a compound. How do we calculate empirical formulas? • We calculate empirical formulas by using the percent composition of each element in a compound and turning it into grams. For example, 68% N becomes 68 g N. We do this for each element. • We then convert the grams of each element into moles using the appropriate conversion and gfm. • Then, we find the lowest number of moles and divide all mole quantities (for all elements) by it in order to get the lowest whole numbers. • If you get a number that is close to a whole number when you divide by the lowest number of moles then round it up to that number. For example, you would round 2.9 to 3. For the number 2.5, you would not round it. You would multiply all mole numbers by 2 so that the 2.5 became 5. • Write the empirical formula using these lowest whole numbers of moles that you found. The number of moles goes next to each element in the compound. Problem • Determine the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen. • Determine the empirical formula of a compound that is 16.49% N, 56.53% O, and 26.97% Na. What are molecular formulas? • The molecular formula is the same as its experimentally determined empirical formula or it is some whole number multiple of it. • To calculate molecular formulas, you first must find the empirical formula of a compound. • Then you must find the gfm of that empirical formula. • Divide the gfm of the molecular formula (usually given to you and called the gram molecular mass) by the gfm of the empirical formula. This should give you a whole number to multiply by. • Multiply the empirical formula by that whole number (that affects the number of each element in the compound) in order to get the molecular formula. Problems • Ethane has a gram molecular mass of 30 g/mol. If ethane’s percent composition is 79.9% carbon and 20.1% hydrogen, what is its molecular formula? What are chemical reactions? • Chemical reactions (sometimes called chemical equations) describe chemical processes. In chemical reactions, an arrow separates the formulas of the reactants (on the left) from the formulas of the products (on the right) • We read reactions from left to right (like a book). • Chemical reactions occur inside of us and in all living things. They help us break down our food intake, send signals throughout our body, and eliminate harmful waste products or toxins. Chemical Reaction (burning of wood) What information can we get from chemical reactions? • We can get the reactants and products of a chemical reaction. We can also get their states. The states of a reactant or product are written next to that reactant or product in subscript parentheses. (s) means solid, (l) means liquid, (g) means gas, and (aq) means aqueous or dissolved in water. • Above the reaction arrow, you might see a chemical formula. That is a catalyst, which helps the reaction go but is not part of the reaction itself. You also could see H, which means that you had to add heat. Problems • Describe the following reactions in words: • CaCO3 (s) CaO (s) + CO2 (g) • S (s) + O2 (g) SO2 (g) What is conservation of mass? • Conservation of mass states that you can neither create nor destroy mass. You can only change its form. Basically, it states that you cannot create something out of nothing. For example, you cannot turn lead into gold. • At the subatomic particle level, (quantum mechanics) we are seeing subatomic particles appear seemingly out of nowhere. This research will give us a better understanding of where these particles came from during and after the initial expansion of the universe (Big Bang) Why is conservation of mass important? • Conservation of mass influences how we write chemical equations. According to conservation of mass, all chemical reactions must be balanced. A balanced chemical reaction has the same number of atoms of each element on both sides of the reaction. How do we balance reactions? • We balance reactions by taking an inventory of each element on both sides of the reaction. • You cannot change the formulas of the reactants or products, you can only change the coefficients out in front of each reactant or product. • Remember that these coefficients affect all elements in the compound and they represent moles. Example of balancing reaction • Balance this reaction: • H2 + O2 H2O • There are two hydrogen on the reactant side and two oxygen on the reactant side. • There are two hydrogen and one oxygen on the product side. This means that the oxygen is unequal. We would balance that by putting a 2 in front of the water. • H2 + O2 2H2O • This balances our oxygen but not hydrogen. We would then have to put what in front of the hydrogen on the reactant side? • Our final equation would look like: • 2H2 + O2 2H2O Practice • Balance these reactions • C2H6 + O2 CO2 + H2O • NaCl + Al(NO3)3 NaNO3 + AlCl3 What are the types of reactions? • There are five types of reactions (know them because I may ask you about them on a quiz…) • Combination reactions occur when two or more elements or compounds combine to form a different compound. A + B AB • Decomposition reactions occur when a compound decays into its component elements or new compounds. AB A +B • Single replacement reactions occur when a metal reacts with an ionic compound. This metal replaces the metal that is in the ionic compound. This forms a new compound and places the other metal on its own. A + BC B + AC where A and B are metals. • Double replacement reactions occur when two ionic compounds react. The metals, in these ionic compounds, switch places and form two new ionic compounds. AB + CD AD + CB • Combustion reactions occur when elements (usually hydrocarbons) burn with oxygen (you must provide an initial jump of energy, like heat, to start them) to produce carbon dioxide and water. CnHn + O2 CO2 + H2O Double replacement reaction Lead Iodide is the solid Combustion in a diesel engine Balance the following reactions and identify their type • C3H8 + O2 CO2 + H2O • H2O2 H2O + O2 • Mg + HCl MgCl2 + H2 What are single replacement reactions? • In a single replacement reaction, one metal reacts with an ionic compound. The metal, that is on its own, replaces the metal that is in the ionic compound. As a result, a new compound forms and a different metal is on its own. How do we write single replacement reactions? • First, you have the metals switch places (if there is only one metal, then the non-metals switch places). Write each element (or polyatomic ion) in its singular form on the product side of the equation. For example: Mg + AlCl3 MgCl + Al • Balance the charges on the new ionic compound that you have formed. Once you have done that, then balance the reaction overall. • Thus it should be: Mg + AlCl3 MgCl2 + Al (balance charges) 3Mg + 2AlCl3 3MgCl2 + 2Al (balance entire reaction) Practice • Write the following reactions: • Mg + Zn(NO3)2 • Al + H2SO4 What are double replacement reactions? • In a double replacement reaction, two ionic compounds react. The metals switch places and form two new ionic compounds. • As a result, two new ionic compounds form. How do we write double replacement reactions? • First, the metals (in both ionic compounds) switch places. Write each element (or polyatomic ion) in its singular form on the product side of the equation. For example: BaCl2 + K2CO3 BaCO3 + KCl • Balance the charges on the new ionic compounds that you have formed. Once you have done that, then balance the reaction overall. BaCl2 + K2CO3 BaCO3 + KCl (charges balanced) BaCl2 + K2CO3 BaCO3 + 2KCl (balance entire reaction) Practice • Write the following double replacement reaction: • MgBr2 + Zn(NO3)2 What are combustion reactions? • In combustion reactions, an element or a compound (usually a hydrocarbon) reacts with oxygen producing heat and light. • Combustion reactions need an initial burst of energy (through a spark) to get going. • Combustion reactions occur when hydrocarbons burn in oxygen to produce carbon dioxide and water. How do we write combustion reactions? • In a combustion reaction, the hydrocarbon burns in oxygen to produce carbon dioxide and water. The products of this type of reaction are always carbon dioxide and water. • Thus we would write: • C3H8 + O2 CO2 + H2O • Then balance the overall reaction to get: • C3H8 + 5O2 3CO2 + 4H2O Practice • Write out and balance the following combustion reactions: • C5H10 + O2 • C6H12O6 + O2 Combustion Reaction What is the activity series of metals? • The activity series of metals (pg. 191 in your textbook) is a list of metals in order of decreasing reactivity. • The activity series determines whether or not a single replacement reaction will occur. Single replacement reactions will only occur if the metal, that is on its own, is more active than the metal that is in the ionic compound. • If the metal that is on its own is not more active than the metal that is in the ionic compound then no reaction will occur. • Remember that the activity series of metals only applies to single replacement reactions. • Write out the following single replacement reactions (write no reaction if a reaction will not occur): • Ag + NaNO3 • Mg + FeCl3 What are net ionic reactions? • You can write single or double replacement reactions as a complete ionic reaction, which is a reaction that shows dissolved ionic compounds as their component ions. You can only split up ionic compounds into their component ions if they have (aq) next to them. You cannot do this for compounds that have (l), (g), or (s). • Spectator ions are ions that are not directly involved in the reaction. These ions appear on both sides of the reaction. You can cancel them out. • After doing this, you can write a net ionic equation that indicates only the particles that actually take part in the reaction. Practice • Write a net ionic equation for the following: AgNO3 (aq) + NaCl (aq) AgCl(s)+ NaNO3 (aq) • You can split the three aqueous compounds into their component ions like so: Ag+ (aq) + NO3- (aq) + Na+(aq) + Cl- (aq) AgCl(s)+ Na+(aq) + NO3- (aq) • Thus, the spectator ions are Na+ and NO3- because they appear on both sides of the reaction unchanged (still in aqueous form). You can cross these out and write the net ionic equation as: Ag+ (aq) + Cl- (aq) AgCl(s) Practice • Write out the following as a net ionic equation: • Pb(s) + 2AgNO3 (aq) Pb(NO3)2 (aq) + 2Ag (s) What is temperature? • Temperature is the measure of the average speed of molecules in a substance. Heat raises the temperature because it causes the molecules to move faster. • Temperature is measured in three scales: Farenheit, Celsius, and Kelvin. Remember that water freezes at 0 C and boils at 100 C. In the Kelvin scale, there are no negative temperatures. • How does temperature affect you? How do you respond to high temperatures? What about low temperatures? Thermometer How do we convert between temperature scales? • There are formulas that allow us to convert between temperature scales. • C=5/9 (F-32) • F=9/5(C) + 32 • K=C + 273 What is absolute zero? • In the Kelvin scale, there are no negative temperatures. 0 Kelvin is called absolute zero because it is the lowest possible temperature. At this temperature, all molecular motion stops, gases occupy no volume, and have no pressure. Temperature Conversions • Complete the following temperature conversions: • Convert 15 C to Farenheit and Kelvin • Convert 98.6 F to Celsius and Kelvin • Convert 373 K to Celsius and Farenheit Temperature Activity • Convert the following temperatures to Kelvin: • 273 C 50 C • 27 C 45 C • 450 C 100 C 75 C 90 C 730 C • Convert the following temperatures to Celsius: • 273 K 0K 50 K • 345 K 400 K 760 K • 540 K 335 K 600 K What is pressure? • Pressure is the force per unit area. Gases move from areas of high pressure to areas of low pressure. • We measure pressure in units of atmospheres (atm), millimeters of mercury (mmHg), Kilopascales (Kpa), and pounds per square inch (psi). Pressure Defined Weight of air on you due to pressure QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are mole ratios? • Mole ratios are a comparison of the number of moles of two different substances (either reactants or products) in a reaction. Mole ratios can compare two reactants, two products, or a reactant and a product. • To find mole ratios, simply look at the coefficients in front of the two substances that you are looking at. Place those two numbers on either side of a colon (:) • For example, in the following reaction: • 3NaNO3 +AlCl33NaCl + Al(NO3)3 • The mole ratio between NaNO3 and Al(NO3)3 is 3 mol NaNO3: 1 mol Al(NO3)3. • What is the mole ratio between AlCl3 and NaCl? What about between NaNO3 and AlCl3? What is Stoichiometry? • Stoichiometry studies the quantitative aspects of chemical reactions. Stoichiometry allows us to predict how much reactant we will need or how much product we will form using those quantities of reactants. What are mole to mole conversions? • In Stoichiometry, we must learn a new conversion called mole to mole conversions. These conversions come from the mole ratios (of coefficients) in a reaction. These conversions allow us to convert from moles of one substance to moles of another substance. • As always, the moles that you want go on top and the moles that you are trying to cancel go on the bottom. Example • 3NaNO3 +AlCl33NaCl + Al(NO3)3 • For the above reaction, how many moles of NaCl can form from 5 moles of AlCl3? • First we must write our given quantity outside of the conversion factor. • Then we must find the mole ratio between NaCl and AlCl3 and it is: • Then we must set up our conversion factor based on the mole quantities from this reaction. The moles that we want (NaCl) go on top and the moles that we want to cancel (AlCl3) go on the bottom of the conversion factor. The numbers of each come from the coefficients in the reaction. You can then convert the numbers. • 3NaNO3 +AlCl33NaCl + Al(NO3)3 • Using the above reaction, how many moles of NaNO3 are needed to react with 2 moles AlCl3? How many moles of Al(NO3)3 can form from 2 moles AlCl3? What are mass to mass conversions? • Mass to mass conversions allow us to convert from grams of one substance to grams of another substance. Mole to mole conversions are still used to convert between moles of two different substances. • If the problem gives you a mass of one reactant or product, you must first convert it into moles. • Then, convert it into moles of a different substance using the mole to mole conversion. • Then, convert from moles of that new substance back into grams of that new substance. In all, you will need up to three conversion factors to do this. How many conversion factors will you need? • You will need 1 conversion factor to go from moles of one substance to moles of another substance. • You will need 2 conversion factors to go from grams of one substance to moles of another substance or from moles of one substance to grams of another substance. • You will need 3 conversion factors to go from grams of one substance to grams of another substance. Practice • 3NaNO3 +AlCl33NaCl + Al(NO3)3 • How many grams of NaCl can form from 3 moles of AlCl3? • How many grams of Al(NO3)3 can form from 30 grams of AlCl3? More Practice • 2Al(NO3)3 +3Zn 3Zn(NO3)2 + 2Al • How many moles of Al can form from 3 moles of Zn? • How many grams of Zn(NO3)2 can form from 60 grams of Zn? What is a limiting reactant? • The limiting reactant is the reactant that is completely used up in a chemical reaction. None of it remains after a reaction becomes complete. It all goes to form products, which is why the limiting reactant determines how much product can form. The limiting reactant is the reactant that forms the least amount of product. • The theoretical yield is how much product (usually in grams) should form in a chemical reaction. The theoretical yield is the amount of product that the limiting reactant forms. The limiting reactant determines how much gas will form How do we determine the limiting reactant? • The limiting reactant is the reactant that forms the least amount of product. You must convert the quantities of reactant to the same product. You can only convert these quantities to the same product because you need to compare them. The reactant that forms the least amount of product is limiting. • Also, the quantity of product that the limiting reactant forms is the theoretical yield. Practice • 2Al(NO3)3 +3Zn 3Zn(NO3)2 + 2Al • If 20 grams Al(NO3)3 and 50 grams of Zn react, which reactant is limiting? What is the theoretical yield for Al? • If 100 grams Al(NO3)3 and 75 grams of Zn react, which reactant is limiting? What is the theoretical yield for Al? • Mg + 2HCl MgCl2 + H2 • Given the above reaction, if 15 grams Mg and 40 grams HCl react, which reactant is limiting? How much Hydrogen (H2) should form? What is the excess reactant? • The excess reactant is the reactant that is left over at the end of a reaction. It is the reactant that forms the larger amount of product when you do limiting reactant calculations. • You can determine how much of the excess reactant is left over at the end of the reaction. • You can tell how much excess reactant is left over by first finding the limiting reactant. • You then can identify which reactant is in excess. • You then take your given quantity of limiting reactant and convert it to grams of the excess reactant in order to see how many grams of excess reactant react with the limiting reactant. • You then subtract the given quantity of excess reactant by how much excess reactant reacts with the limiting reactant. This tells you how much excess reactant is left over at the end of the reaction. Practice • Mg + 2HCl MgCl2 + H2 • Given the above reaction, if 15 grams Mg and 40 grams HCl react, which reactant is in excess? How much of this reactant is left at the end of the reaction? What is the percent yield? • Percent yield compares how much product you made (actual yield) to how much product you should have made (theoretical yield). A percent yield above 90% is very good. • Percent yield=(actual yield/theoretical yield) x 100 Practice • Mg + 2HCl MgCl2 + H2 • Given the above reaction, if 1 gram of H2 forms, what is the percent yield? Saturation Example QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Example of how pressure affects gas solutes. • Increasing the pressure above the solution makes more gas molecules dissolve in a solvent. • What companies would use this principle? QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What is concentration? • Concentration is the amount of solute that is dissolved in a given amount of solvent. Solutions are concentrated if they have a large amount of solute dissolved in a given amount of solvent. • Why would we care about the concentration of a substance? • All substances are toxic to humans in certain concentrations. Concentrated Detergent What do we measure concentration in? • We measure concentration using four different measurements: • Molarity (M)=moles solute/liters solvent • Grams per Liter (g/L)=grams solute/liters solvent • Which is easier to determine? Practice • I dissolve 60 grams of Pb(NO3)2 into .5 L of water. What is the concentration of this mixture in molarity? What about grams per Liter? • I dissolve 40 grams of NaOH into 2 L of water. Express this concentration in molarity and grams per Liter. What are the other concentration units? • Parts per million (ppm)= milligrams of solute/kilograms of solvent • Percent Composition (%)= (grams solute/mL of solvent)x100 Practice Problems • A sample of river water has 4 mg of lead in .2 Kg of water. Express this concentration in ppm. • A fertilizer says that it has “5% NH4NO3 by volume.” You test this batch of fertilizer and find that it has 10 mg NH4NO3 per 500 mL bottle. Is this bottle’s label true in terms of how much NH4NO3 it actually has? What is heat? • Heat is a form of energy that flows from higher temperatures (hot) to cold (lower temperatures). Heat is released or absorbed during any mechanical work. • You cannot feel heat, you only feel heat transfer. • Exothermic-heat is released. Endothermic-heat is absorbed. • Why do you feel cold on a winter day? Image of heat transfer QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What is a phase change? • The phases of matter are solid, liquid, and gas. A phase change occurs when substances change their state of matter. This is a physical change. • During a phase change, the temperature of a substance does not change. All of the absorbed energy is used to make the substance change state. What are the phase changes? • Solid to liquid: Melting • Liquid to Solid: Freezing • Liquid to gas: Boiling/Evaporation • Gas to liquid: Condensation • Solid to gas: Sublimation (see picture) • Gas to Solid: Solidication. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. • System- This is the specific part of the universe that you are focusing on. • Surroundings- This is everything outside of the system. • Endothermic reaction-This is a process that absorbs heat from the surroundings. • Exothermic reaction-This is a process that releases heat to the surroundings. What is specific heat? • Specific heat is the amount of heat needed to heat 1 g of a substance by 1 C. Every substance has its own unique specific heat. • We can use this formula if we have specific heat: H=cm T, where H equals the heat change in joules, c=specific heat, m=mass, and T=Tfinal-Tinitial. Heat calculations • How much heat is needed to raise the temperature of water from 30 C to 50 C? (Specific heat of water=4.18 J/g C • How much heat is released when 10 g copper cools from 100 C to 50 C? Specific heat of copper=.33 J/g C How do we use ∆H=cm∆T for two substances? • In this situation, one substance loses heat (the one at the higher temperature) and one substance absorbs that same amount of heat (the one at lower temperature). Since they absorb or release the same amount of heat, we can relate them. Also, the final temperature of the system is the same for both substances. • We use this equation: -cm∆T = cm∆T Heat released heat absorbed Where ∆T=Tfinal - Tinitial • We use this formula to find specific heats of objects among other things. • I drop a 20 g piece of copper, at 100 C, into 20 g of water at 25 C. What is the final temperature of my system? Cwater=4.18 J/g C and Ccopper=.387 J/g C • I drop 20 grams of an unknown substance, at 90C, into 50 grams of water at 20 C. If the final temperature of the water and unknown substance are 24 C, what is the specific heat of the uknown substance? How do we calculate heat for phase changes? • For phase changes, we use the following formulas: • H=m Hfus H=m Hvap • Hfus=heat it takes to melt or freeze something. • Hvap=heat it takes to boil or condense something. Image of Ice Melting QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Heat calculations • How much heat does it take to melt 20 grams of ice at 0 C? (Heat of fusion=333 J/g) • I add 400 J of heat to 40 g of ice. How much of this ice melts? How much of it is left over? How can we use the additive properties of heat? • We can use the additive properties of heat in problems that involve both phase and temperature changes (for example, a problem that goes from ice at one temperature to water at a different temperature). When a temperature change occurs, use cm∆T but use either m∆Hfus or m∆Hvap for phase changes. • Add up the heats that you calculate as you change the temperature. Problems • How much heat is needed to melt 50 g ice, at -10 C into water at 15 C? • How much heat is needed to change 20 g of water, at 50 C, into gas at 110C? What is Pressure? • Pressure is a force per unit area. Pressure is the force per unit area. Gases move from areas of high pressure to areas of low pressure. • We measure pressure in units of atmospheres (atm), millimeters of mercury (mmHg), Kilopascales (Kpa), and pounds per square inch (psi). • We can convert between pressure units. How do we convert between pressure units? • Like any conversion, the units you want go on top of the conversion and the units that you want to cancel go on the bottom of the conversion. • Use these equivalencies: 1 atm= 760 mmHg = 101.3 kPa Problems • Convert 5 atm into kilopascales. • What represents more pressure, .25 atm or 550 mmHg? What is Charles’ Law? • Charles’ law is the volume-temperature relationship. It is a direct relationship and pressure and amount of gas are kept constant. • V1/T1=V2/T2 • V1-initial volume, T1=initial temperature, V2-final volume, T2-final temperature. • Remember that temperatures must be in Kelvin!! Charles’ Law Example Calculations • I have a gas at 30 K and 4 L. What is my new volume when my temperature increases to 60 K? • I have a gas at 25 C and 2 L. What is my new temperature if the volume decreases to .5 L? What is Boyle’s Law? • Boyle’s law is the pressure-volume relationship. It is an indirect relationship and temperature and amount of gas are kept constant. • P1V1=P2V2 • V1-initial volume, P1=initial pressure, V2-final volume, P2-final pressure. Too Much Pressure… Calculations • I have a gas at 2 atm and 3 L. What is my new pressure if my volume increases to 6 L? • I have a gas at 400 kPa and 7 L. What is my new volume if the pressure increases to 500 kPa? What is Gay-Lussac’s Law? • Gay Lussac’s law is the pressure-temperature relationship. It is a direct relationship and volume and amount of gas are kept constant. • P1/T1=P2/T2 • P1-initial pressure, T1=initial temperature, P2-final pressure, T2-final temperature. • Remember that temperatures must be in Kelvin!! Example Calculations • I have a gas at 600 mmHg and 30 C. What is my new temperature if my pressure increases to 800 mmHg? • I have a gas at 400 mmHg and 20 C. What is my new pressure if my temperature increases to 100 C? What gas laws have we learned so far? • Charles: V1/T1=V2/T2 (Pressure is constant) • Boyles: P1V1=P2V2 (Temperature is constant) • Gay-Lussac’s: P1/T1=P2/T2 (Volume is constant) • Remember to use Kelvin and that the amount of gas is constant in all of these. • I have a gas at 3 L and 5 atm. What is my new pressure if the volume decreases to .5 L? • I have a tire at a pressure of 32 psi and a temperature of 30 C. What is my new temperature if the pressure becomes 45 psi? What is the Combined Gas Law? • The combined gas law is a combination of the other three laws that we have learned. You can derive the other three gas laws by keeping one of the variables, in the combined gas law, constant. • P1V1/T1=P2V2/T2 • Only the amount of gas is kept constant. P1-initial pressure, V1-initial volume, T1=initial temperature, P2-final pressure, V2-final volume, T2-final temperature. Calculations • A sample of air occupies a volume of 3 L at 50 C and 4 atm. What is the volume of air at 100 C and 8 atm? • A sample of air occupies a volume of 4 L at 30 C and 2 atm. What is the new pressure of this gas at 2 L and 25 C? What are acids? What are bases? • Acids are hydrogen donors while bases are hydrogen acceptors. • Acids are corrosive and taste sour. Bases are bitter, slippery, and caustic. Both can cause serious chemical burns. • Strong acids completely dissociate in water while weak acids only partially dissociate in water. The same is true for strong bases vs. weak bases. • What are some examples of acids or bases? Strong vs. Weak acids • These images show the difference between strong and weak acids. • Does the same thing hold true for strong bases vs. weak bases? Quic kTime™ and a TIFF (Unc ompres sed) dec ompres sor are needed to see this pic ture. Quic kTime™ and a TIFF (Unc ompres sed) dec ompres sor are needed to see this pic ture. Strong Acids can be dangerous What is the pH scale? • The pH scale is a scale that goes from 0 to 14. • The pH scale allows us to measure acidity and basicity. We classify substances according to their pH. pH is a measure of the hydrogen ion concentration. pH=-log[H+] pOH=-log[OH-] pH+pOH=14 Here is the pH scale QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. pH classifications • Please label the following classifications on your pH scale: • pH 0-3=strong acid • pH 4-6=weak acid • pH6-8=neutral • pH 9-11=weak base • pH 11-14=strong base • If something has a pH of 5, what do we classify it as? What are reaction rates? • Reaction rates measure how fast a reaction occurs. They are proportional to the decrease in reactant concentration over time or the increase in product concentration over time. • Consider this reaction: AB • Rate= -[A]/ t= [B]/ t What affects reaction rates? • In order for a reaction to occur, the reactants must have enough energy to overcome the activation energy barrier (see board). The reaction, even if it is spontaneous, will not occur unless it can overcome the activation energy barrier. • Temperature: This affects reaction rate because temperature measures the speed of molecules. If the temperature is high then the molecules are moving faster and you will have more collisions with higher energy. This will increase the reaction rate. At a lower temperature, the molecules move slower and you will have less high energy collisions, which will slow down the reaction rate. • Pressure affects reaction rates in the same way that temperature does. • Concentration-this is the amount of each reactant. If you increase the concentrations of the reactants, there are more reactants in the vessel. This increases the probability of high energy collisions and increases the reaction rate. Decreasing the concentrations of reactants decreases the reaction rate because there are less reactants and less high energy collisions. • Surface Area- increasing the surface area increases the reaction rate because there are more high energy collisions between reactants. What is a catalyst? • A catalyst is something that increases the reaction rate by lowering the activation energy. A catalyst can be a metal ion, or a huge protein. The biological catalysts are enzymes. Example of a Catalyst • This is a picture of an enzyme in your body. • You can see where it binds the reactants so that they can react. This binding lowers the activation energy. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What is Equilibrium? • Equilibrium reactions are represented by a double arrow. For example: AB. You cannot apply normal stoichiometric calculations to these reactions because there is both a forward and reverse reaction that can occur. • Equilibrium occurs when the forward and reverse reaction rates are equal. What is Le-Chatlier’s Principle? • Le-Chatlier’s principle states that if a stress is applied to a system at equilibrium then the system changes to relieve that stress. • This shift will either be toward the product side or toward the reactant side depending on the stress that is put on the equilibrium. This only applies to aqueous and gaseous reactants or products. What affects equilibrium? • Concentration-Increasing the concentration of a product pushes the equilibrium reaction back towards the reactant side. Increasing the concentration of a reactant pushes the equilibrium reaction in the direction of the product side. • Decreasing the concentration of a product pulls the equilibrium towards the product side. Decreasing the concentration of a reactant pulls the equilibrium towards the reactant side. Temperature • Increasing the temperature of an equilibrium causes the reaction to shift towards the side that is endothermic (absorbs heat) or does not have the heat term. • Decreasing the temperature of an equilibrium causes the reaction to shift towards the side that is exothermic (releases heat) or the side that has the heat term. Pressure • Pressure only affects gaseous reactants or products (g) not aqueous reactants or products. It only affects equilibrium systems with an unequal number of moles of gaseous reactants and products. • Increasing the pressure forces the equilibrium to shift towards the side with fewer moles of gas. • Decreasing the pressure forces the equilibrium to shift towards the side with more moles of gas. Problem • Consider the following equilibrium reaction: N2(g) + 3H2(g) 2NH3(g) + heat Which way does the equilibrium shift if I make the following changes to the system: Increase N2 Decrease H2 concentration Decrease NH3 concentration Increase temperature Decrease temperature Increase pressure Example of Le-Chatlier’s Principle • Ammomia Synthesis (the Haber process): N2(g) + 3H2(g) + heat 2NH3(g) • How can I maximize the synthesis of ammonia using LeChatlier’s principle? QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Atomic Structure Review • What did we learn about in our atomic structure unit? • Atoms-definition, number of particles in each, isotopes, isotope notation. • Periodic table families and trends. • Nuclear Chemistry-types of nuclear decay, fission, fusion, and writing nuclear decay reactions. • Electron configurations-long-hand and short-hand along with finding the number of valence electrons. Problems • How many protons, neutrons, and electrons are in carbon 14? • Which has the higher electronegativity between C, N, and O? • Write out the nuclear decay reaction that occurs when U-235 undergoes alpha decay. • What is the electron configuration for bromine and how many valence electrons does it have? What did we learn about bonding? • We learned about the types of bonding-ionic, metallic, covalent, and covalent network. • We learned about how ions form and how to make ionic compounds. • We learned how to name compounds. • We learned how to make Lewis structures to represent covalent structures. • We learned about some biochem topics also. • What type of bonding do we find in CO2? • Form an ionic compound with Mg and Cl, and Mg and NO3. • Name those compounds from second bullet. • Draw the structure of CH4, N2H4. • What are the subunits of proteins? What did we learn about moles and reactions? • We learned how to convert moles to mass, volume at STP, and particles. We learned two step conversions, gfm, empirical formulae, and molecular formulae. • We learned conservation of mass and how to balance reactions. • We learned how to write reactions. Problems • Convert 22 grams of Mg(NO3)2 to moles. • Convert 20 L of oxygen gas to moles • Convert 3x1024 atoms of nitrogen into grams of nitrogen. • Balance the following reaction: • NaCl + Al(NO3)3 NaNO3 + AlCl3 What did we learn about Stoichiometry? • We learned how to do mole to mole conversions, mass to mass conversions, calculate limiting reactants, and how to calculate excess reactants. • 2Al(NO3)3 +3Zn 3Zn(NO3)2 + 2Al • If 20 grams Al(NO3)3 and 50 grams of Zn react, which reactant is limiting? What is the theoretical yield for Al? • What did we learn about solutions? • We learned about how substances dissolve, what makes them dissolve, how we can get more solute to dissolve into a solvent, what a solute and solvent are, and how to calculate concentration. • How can I get more solute to dissolve into a solvent? • What is the molarity of a solution where I dissolve 20 g of MgCl2 into 2L of water? What did we learn about acids/bases, reaction rates, and equilibrium? • The definition of an acid/base, the difference between strong acids/bases and weak acids/bases, the pH scale and classifications. • For reaction rates, we learned the definition and what we can do to speed up a reaction rate. • For equilibrium, we learned the definition and how to apply Le-Chatlier’s principle. • What is an acid? How would we classify a solution that has a pH of 9? • How can we increase the rate of a reaction? • Consider the following equilibrium reaction: N2(g) + 3H2(g) + heat 2NH3(g) Which way does the equilibrium shift if I make the following changes to the system: Increase N2 Decrease H2 concentration Decrease NH3 concentration Increase temperature Decrease temperature Increase pressure What did we learn about thermodynamics? • We learned about the types of phase changes, the definition of heat, how to calculate heat during a temperature change, or phase change, or both. • How much heat is needed to raise 20 g of water’s temperature from 30 C to 50 C? (Specific heat of water=4.18 J/g C • How much heat is needed to melt 20 grams of ice? Heat of fusion for ice=333 J/g What did we learn about the gas laws? • We learned the definition of temperature, pressure, absolute zero, and how to use the four gas laws. • A sample of air occupies a volume of 3 L at 50 C and 4 atm. What is the volume of air at 100 C and 8 atm? • I have a gas at 2 atm and 3 L. What is my new pressure if my volume increases to 6 L? Warm-up • Why would we want to know how much energy it takes to melt ice? Why is this important to us? QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Exit Slip • What improvements could you make to the experiment in order to improve your results? QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What is Hess’ Law? • Hess’s Law of Heat Summation states that if you add two or more thermochemical equations to give a final equation, then you can also add the heat changes to give the final heat change. • In order to use Hess’s Law, you must cancel out substances or elements (if they are on their own) that are on both sides of the reaction when you add the reactions. • When you use Hess’s Law, make sure you identify which reaction has the reactants or products that you want. • If the product you want is written as a reactant, then you can switch the reaction around. If you do this, the sign on the heat (∆H) changes (it can go from a negative to a positive). • If you need to multiply a reaction (in order to decrease or increase the amount of a product or reactant) by a constant, then you must multiply ∆H by that same constant. • Once you have the needed reaction, add up the heats. Solve a Hess’s Law problem • Solve this Hess’s law problem: Find ∆H for the following reaction: 2Cu(s) + O2(g)2CuO(s) Given: CuO(s) + Cu(s) Cu2O(s) ∆H=-11.3 KJ Cu2O(s) + 1/2O2(g)2CuO(s) ∆H=-114.6 kJ Which reactions contain the reactants you want? Which ones have the products that you want? What is the Ideal Gas Law? • The ideal gas law relates temperature, pressure, and volume to amount of gas in moles. • This law forces us to make ideal gas assumptions which are: a. There are no attractive forces between gas molecules. b. Gas molecules occupy no volume. • Real gases do not behave this way all the time, but the ideal gas law gives us a good approximation. • The ideal gas law is this formula: PV=nRT P=pressure in atm, V=volume in L, n=moles of gas, R=.0821, T=Temperature in Kelvin. • You can make changes to this equation in order to make it work better for real gases. Example Problems • I have a gas at 2 atm of pressure, 2L of volume, and 30 C. How many moles of this gas do I have? • 50 grams of Carbon dioxide gas is at a temperature of 150 C, and a volume of 3 L. What is the pressure of this gas? • I have oxygen gas in a container at 3 atm, 5 L, and 25 C. How many grams of oxygen gas are in this container? What is Dalton’s Law of Partial Pressures? • Partial Pressure is the pressure contribution that each gas, in a mixture, makes to the total pressure of the gas mixture. • Dalton’s law of partial pressure states that at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases. • Ptotal=P1 + P2 + P3 +… Example Problem • Air contains oxygen (O2), nitrogen (N2), carbon dioxide (CO2) and trace amounts of other gases. What is the partial pressure of oxygen P(O2) at 101.3 kPa of pressure if P(N2)=79.1 kPa, P(CO2)=.04 kPA, and P(other gases)=.94 kPA? What are oxidation and reduction reactions? • Oxidation and Reduction reactions (called redox) occur when electrons are transferred between reactants. What is oxidation? • Oxidation occurs when an organic compound adds oxygen or decreases its number of hydrogens. • We also define it as the loss of electrons. • Substances that oxidize will donate electrons, which means that they are a reducing agent. What is reduction? • Reduction is when organic compounds lose oxygen or gain hydrogens. • We also define it as the gain of electrons. • The substance that is reduced in an oxidation-reduction reaction is called the oxidizing agent. Redox Reaction Example What are the rules for assigning oxidation numbers? 1. 2. 3. 4. The oxidation number of a monoatomic ion is equal to its ionic charge. (Sodium has a +1 charge) The oxidation number of hydrogen in a compound is always +1 except in metal hydrides such as NaH (where it is -1) The oxidation number of oxygen in a compound is always -2. The oxidation number of any uncombined element is zero. For example, the oxidation number of solid potassium metal (K) is zero. What about the oxygen in O2? 5. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal zero. 6. For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion. • What is the oxidation number of each element in the following compounds? a. SO2 b. CO3-2 c. K2SO4 How do we identify oxidation and reduction in reactions? • An increase in oxidation number signifies oxidation. • A decrease in oxidation number signifies reduction. • In order to determine which element is oxidized or reduced we must assign oxidation numbers to each element, on both sides of the reaction. • Then we determine which element had a decrease in its oxidation number (reduced) and which element had an increase in its oxidation number (oxidation). • Use the information on the previous slide to solve the following problem: Use the changes in oxidation numbers to identify which elements are oxidized and reduced in this reaction: Zn(s) + 2MnO2(s) + 2NH4Cl(aq) ZnCl2(aq) + Mn2O3(s) + 2NH3(g) + H2O(l) How do we balance redox reactions? • We use the oxidation number change method to balance redox reactions by comparing the increases and decreases in oxidation numbers. • This principle works because the total number of electrons gained in reduction must equal the total number of electrons lost in oxidation. • In order to balance redox reactions we must use the following steps: 1. 2. 3. 4. 5. Assign oxidation numbers to all the atoms in the equation. Identify which atoms are oxidized and which are reduced. Use a line to connect the atoms that undergo oxidation and those that undergo reduction. Write the oxidation number change at the midpoint of each line. Make the total increase in oxidation number equal to the total decrease in oxidation number by using the appropriate coefficients. Finally, check to be sure that the equation is balanced for both atoms and charge. Use the method highlighted above to solve the following problem: • Balance this redox reaction by using the oxidation number change method. K2Cr2O7(aq) + H2O(l) + S(s)KOH(aq) + Cr2O3(aq) + SO2(g) How do we use half reactions to balance Redox reactions? • Half-reactions are equations that show either the reduction or the oxidation of a species in an oxidation-reduction reaction. • The half-reaction method is used to balance redox reactions by balancing the oxidation and reduction half reactions. • This method is particularly useful when you are balancing ionic reactions. What are the steps to balancing redox reactions using half reactions? 1. 2. 3. 4. 5. Write the equation in ionic form. Identify the oxidized species and the reduced species. Write out their half reactions. Balance each half reaction in terms of charge. Make sure that the charges are equal on both sides of the half reaction. Multiply each half reaction by an appropriate number to make the electron charges equal. This will allow you to cancel out electrons in each half reaction. Combine half-reactions. Check to make sure all numbers are balanced. Use the method outlined above for the following reaction: • Balance this reaction using the half-reaction method: MnO4-(aq) + Cl-(g)Mn+2(aq) + Cl2(g) But this reaction does not balance!! • The above reaction is not balanced for the oxygen atoms. How can we balance it? • Under acidic conditions, we do the following: • Add oxygen to the either the reactant or product side using H2O. Add as many waters as you need to balance out the oxygen. (If you have four oxygens on the reactant side, you will need four waters on the product side). Continuing under Acidic conditions: • Add H+ to the opposite side so that we can balance out the hydrogens (if you added four waters to one side, then you would add eight H+ ions to the other side) that we just added to the other side of the reaction when we added water. • Check if charges balance. What about under basic conditions? • Under Basic conditions, we balance the reaction under acidic conditions then we add OH- to both sides of the equation. The number of hydroxides (OH-) added equals the number of H+ ions. • Together, the H+ and OH- ions form water. Use these new waters to cancel out the waters on the other side of the reaction. Now back to our problem: • Balance this reaction under acidic and basic conditions: MnO4-(aq) + Cl-(g)Mn+2(aq) + Cl2(g) Electrochemistry • Electrochemical cells convert chemical energy into electrical energy and vice versa. All electrochemical processes involve redox reactions. • An electric current is produced from the flow of electrons between the ions in the redox reaction. Volataic Cells • Voltaic cells are electrochemical cells that are used to convert chemical energy into electrical energy. The energy is produced by spontaneous redox reactions within the cell. • In order to create a voltaic cell, you must split the reaction into two half reactions. A half-cell is one part of a voltaic cell in which oxidation or reduction occurs. The two halfcells are separated by a porous partition. • A salt bridge is a tube that contains a conducting solution. A salt bride allows the passage of ions from one half cell to the other it prevents the two solutions from mixing completely. How do voltaic cells work? • The half-cell (or electrode) where oxidation occurs is called the anode. • The half-cell where reduction occurs is called the cathode. • Electrons flow from anode to cathode. If you connect the two half-cells via an external wire (and with a salt bridge) you can generate an electrical current, which is how batteries work. Here is an example of an electrochemical cell: How do we calculate cell potentials? • In order to calculate cell potentials, we must look at the standard reduction potential. • For half reactions, the standard reduction potential is always written showing the species undergoing reduction. Electrons are reactants. • For standard reduction potential you must have standard conditions (25 C, and 1 atm of pressure). • The standard reduction potential for reducing H+(aq) to H2(g) is defined to be 0 V • In order to calculate standard potential (E) for any electrochemical cell we must use this formula: • E= Eanode+ Ecathode • The anode is the reaction with the more negative half-cell reduction potential. • The cathode is the reaction with the more positive half-cell reduction potential. • For the anode, you will change the sign on its voltage because that voltage is for reduction. To change the voltage to oxidation, you must change the sign on the E for that reaction. • For any electrochemcial cell to be spontaneous (occur without the input of energy) its standard cell potential must be positive. • What is the standard cell potential (E) for the following redox reaction and is this reaction spontaneous? Zn(s) + Cu+2(aq)Zn+2(aq) + Cu(s) Make sure you identify the oxidation and reduction reactions. Look up E for each half reaction (remember to change sign on voltage for anode) in your textbook. What are alkanes? • Alkanes are hydrocarbons that contain only single covalent bonds. They can be in a long and straight chain molecule or they can be branched. • We name the straight-chain alkane based on its number of carbons. We use the following prefixes and add -ane to the end: What are functional groups? • Functional groups are a specific arrangement of atoms in an organic compound that is capable of characteristic chemical reactions. • When we represent the general form of functional groups, we use an -R group to represent carbon chains that are attached to the functional group. • The table on pg. 712 (AW book) shows the general compound forms and their functional groups. This table should be on your note page for the upcoming organic chemistry quiz… Prefixes based on # of carbons in a straight chain: 1 carbon-meth 2 carbons-eth 3 carbons-prop 4 carbons-but 5 carbons-pent 6 carbons-hex 7 carbons-hep 8 carbons-oct 9 carbons-non 10 carbons-dec Images of alkanes QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. How do we name alkanes? 1. 2. 3. 4. 5. Find the longest continuous chain of carbons in the molecule. If it is four carbons, then you call it butane (as the parent structure) Number the carbons in the main chain in sequence. This will allow you to identify each functional group’s position on the chain. Add numbers to the names of the substituent or functional groups to identify their positions on the chain. Use prefixes to indicate the appearance of a group more than once in the structure. List the names of the functional groups in alphabetical order. How do we name the functional groups? • If the functional group is a hydrocarbon then it is called an alkyl group. You name them by removing the -ane end and adding -yl to the end. Make sure you number where they are on the chain when you name the compound. • For halogens, the are named by taking off -ine at the end and adding -o (chlorine becomes chloro). Make sure you number where they are on the chain when you name the compound. • Halocarbons are synthesized and used as anesthetics and insecticides. How do we name alkenes and alkynes? • Alkenes are organic compounds containing carbon-carbon double bonds. • Alkynes are organic compounds containing carbon-carbon triple bonds. • Organic compounds that contain double and triple bonds are called unsaturated compounds because they contain fewer than the maximum number of hydrogens in their structure. • Alkanes that contain the maximum number of hydrogens are called saturated compounds (no carbon-carbon multiple bonds). • To name alkenes or alkynes, follow the same naming rules and numbering. However, the continuous chain carbon must contain the double or triple bond. The ending changes from -ane to ene for alkenes or -yne for alkynes. • In naming these compounds, you must number the chain so that the carbon atoms of the double or triple bond gets the lowest possible numbers. Make sure to put the carbon number where the double bond is on the alkene or alkyne when you name the compound. Name the compounds on the board What are the types of organic reactions? • Substitution: The replacement of an atom or group of atoms by another atom or group of atoms. For example, a halogen can replace the hydrogen on an alkane to produce a halocarbon. • Let’s do some examples on the board. Many reactions we see involve double or triple bonds • Addition reaction-this occurs when a substance is added at the double or triple bond of an alkene or alkyne. Addition reactions allow us to introduce new functional groups into organic molecules. For example, you can add two chlorines on the carbons that make up the double bond. • If you add water to an alkene then it is called a hydration reaction. • Hydrogenation reaction-this occurs when you add hydrogen (H2) to a carbon-carbon double bond and produce an alkane. This reaction requires a catalyst in the form of platinum (Pt) or palladium (Pd) What about oxidation/reduction? • Dehydrogenation reactions occur when you lose hydrogen and create a carbon-carbon double or triple bond. Strong heating and a catalyst are necessary to do this. This causes oxidation of the molecule (loss of hydrogen is considered oxidation). • You can oxidize molecules by adding more oxygens. Alkanes can become alcohols if you oxidize them. These alcohols can become aldehydes or ketones after oxidation depending on their functional groups. These can then be oxidized into esters or alcohols depending on their functional groups. • Hydrogenation reactions are reduction (adding more hydrogen) What is hydrolization? • Hydrolization of esters occur if an ester is heated with water for several hours in strong acid or base solutions. This produces a carboxylic acid and an alcohol. What is Biochemistry? • Biochemistry covers large macromolecules that are found in most life forms. • These macromolecules (large molecules) form from polymers, which are continuous repetitions of the same smaller molecules (monomers). • The types of macromolecules are: carbohydrates, lipids (fats), proteins, and nucleic acids (DNA). What are carbohydrates? • Carbohydrates are monomers and polymers of aldehydes and ketones that have numerous hydroxy groups attached. They have the formula of Cn(H2O)n. • The simplest carbohydrates are simple sugars called monosaccharides such as glucose and fructose (fruit sugar). • If two sugars come together via a polysaccharide bond, then we call them disaccharides. The most famous is a combination of fructose and glucose, which creates sucrose. This is the sugar that we use in candy and to sweeten our drinks. Image of simple sugars QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are polysaccharides? • The linkage of many monosaccharide monomers produces polysaccharides. • Examples include: Starch-it is made up of glucose molecules and serves to store glucose molecules in plants. Starch can partially dissolve in water and is edible. Glycogen-this is the animal form of starch. Humans store glucose in this way Cellulose-it is made up of glucose. The orientation of the bond linking the glucose monomers is different from the orientation in starch. As a result, cellulose is insoluble and can only be digested by a limited number of organisms. Continuation of cellulose • Cellulose forms a rigid structure with other cellulose molecules. Cellulose is used as a structural polysaccharide. Plant cell walls, as in wood, are made of cellulose. • Humans cannot digest cellulose. Cellulose is then considered what type of nutrient? Image of Cellulose QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Lipids • Lipids are a large class of relatively water-insoluble compounds that include fats, oils, and waxes. They can dissolve in organic solvents such as ether. • Triglycerides are triesters of long-chain fatty acids and glycerol. They are the major component of animal fats and oils. Triglycerides are simple lipids and they are important as the storage form of fat in the human body. • Triglycerides are essentially esters that we can hydrolize by boiling them with aqueous sodium hydroxide. This makes soap. What are phospholipids? • Phospholipids are a special type of lipid that contain a polar head (phosphate group) and a non-polar tail (hydrocarbons). • These unique lipids make up the cell membrane in animal cells. It is called the phospholipid bilayer. Image of a Phospholipid QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Image of a phospholipid bilayer QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. • • The polar head faces the outside and inside of the cell so that it can interact with molecules outside or inside of the cell. The non-polar tails face toward the inside on both layers so that they can interact with one another. This makes the cell membrane relatively impermeable because molecules have a hard time getting through the nonpolar tails. Large molecules cannot fit between the phospholipids. Here are some images of Lipids QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What is the difference between saturated vs. unsaturated fats? • Saturated fats have no carbon-carbon double bonds in their structure. This allows them to interact well (VDW forces) and melt at higher temperatures. • Unsaturated fats have carbon-carbon double bonds in their structures. This puts a kink in the fatty acid chains, which means that they do not interact as well (VDW forces) and melt at lower temperatures. They are usually liquid at room temperature. • What household products are made up of mostly saturated fat? Which ones are made up of most unsaturated fats? Unsaturated vs. Saturated Fats • Here are two images that show the difference between saturated and unsaturated fats. QuickTime™ and a TIF F (Uncompressed) decompressor are needed to see this picture. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are amino acids? • Amino acids are compounds that contain amino (NH2) and carboxylic acid (-COOH) groups in the same molecule. • There are 20 amino acids, all with the aforementioned functional groups. They also have another functional group and hydrogen around the central carbon. • At human blood pH (7.4), these amino acids exist in what is called their zwitterion form. Here is an image of amino acid and zwitterion form (know it) QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are peptides? • Amino acids are the building blocks of proteins. They come together via peptide bonds, which form between one carboxylic acid of one amino acid and the amino group of the next amino acid. • Peptides are any combination of amino acids that have come together via peptide bonds. • Polypeptides are peptides with more than 10 amino acids. • Proteins are peptides with more than about 100 amino acids. Here is an image of a peptide bond QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. How do proteins get their various shapes? • Proteins fold into various shapes known as alpha helices or beta sheets. They may fold into more complex shapes based on intermolecular forces, covalent bonds formed between cysteine molecules, and electrostatic attractions. • Proteins can be denatured by changes in temperature and pH. These changes disrupt the attractions that keep them folded in the correct conformation. Images of protein shapes • Here is an image of an alpha helix. • Here is an image of a beta sheet. QuickTi me™ and a TIFF ( Uncompressed) decompressor are needed to see thi s pi ctur e. QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are Enzymes? • Enzymes are proteins that act as biological catalysts. Enzymes are unchanged during a chemical reaction. Enzymes help necessary body reactions to occur. • The enzyme binds substrates (the molecules on which an enzyme acts) at its active site. This is where the enzyme converts substrates into products. The enzyme holds the substrates in place (through intermolecular forces) so that they can react or it even reacts with the substrate. Each enzyme is specific to a certain substrate. What affects enzyme function? • Some enzymes will not function without a coenzyme. These are metal ions or small organic molecules that need to be present in order for an enzyme catalyzed reaction to occur. Many coenzymes are nutrients that we take in everyday such as zinc, iron, and ascorbic acid (vitamin C). • Enzymes can only function in certain temperature or pH ranges. They break down and do not function if they are not within that range. What body conditions (due to sickness) can impair enzyme function? • Some poisons inhibit enzyme function. They attach to the enzyme at the active site and prevent the substrate from binding. They can also attach elsewhere on the enzyme and change its shape so that it cannot bind the substrate. Here is an enzyme: QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. What are nucleic acids? • Nucleic acids are polymers found primarily in cell nuclei. The most common are DNA (deoxyribonucleic acid) or RNA (ribonucleic acid). DNA stores the information needed to make proteins and governs the reproduction and growth of cells and new organisms. RNA transmits the information stored in DNA. • Nucleotides are the subunits of DNA and RNA. They contain a phosphate group, a five carbon sugar (ribose), and a nitrogen base. For DNA its ribose is missing an oxygen. The four bases are adenine (A), Thymine (T), Cytosine (C), and Guanine (G). How does DNA come together? • DNA winds in a double helix shape where the negative charges on the phosphate groups repel each other (this keeps the helices at the correct distance apart). In the middle, the bases interact through hydrogen bonding. • The bases pair in this way:A-T and C-G. RNA contains these bases but it does not have Thymine, it contains Uracil (U), which pairs up with adenine. • Find the corresponding DNA sequence: TTCGATTGGCAAA Here is the DNA double helix: QuickTime™ and a TIFF (Uncompressed) decompressor are needed to see this picture. Final Exam Review • The following items will be on our semester 2 final: • Stoichiometry, solutions, acids and bases, reaction rates, equilibrium, thermodynamics (including Hess’s law), gas laws (including Ideal Gas Law), oxidation/reduction, organic chemistry, and biochemistry. Review Problems • Find the oxidation numbers on the elements in the following compounds: CaSO4 MgO O2 • In the reaction below, determine the oxidized element, reduced element, oxidizing agent, and reducing agent. Mg + CaSO4 MgSO4 + Ca • What is the pressure of 3 mol of gas at a volume of 3 L and a temperature of 30 C? O. Chem + Biochem Problems • What is the complimentary DNA sequence to AATCCG? • What molecule makes up a cell membrane? What does it form? • What are the subunits of proteins? What is the zwitterion form of this molecule? What is an enzyme? What affects its activity? • What are the polysaccharides? • Name the compounds on the board (use IUPAC rules) and write out the reactions that are there.