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Lyon Midterm Review Packet You are responsible for all the material we’ve covered this year. The review material below, and the practice problems that follow, does not cover everything we’ve done so far (such as empirical formulas). Things below we haven’t covered yet (like Gas Laws) will not be on the midterm: Questions in the midterm will not necessarily be addressed in this packet. Measurement Quantity Being Measured Units of Measurement Unit Symbols mass kilogram, gram, centigram, milligram kg, g, cg, mg volume kiloliter, liter, milliliter, cubic meter, cubic decimeter, cubic centimeter, cubic millimeter kL, L, mL, m3, dm3, cm3, mm3, length kilometer, meter, centimeter, millimeter km, m, cm, mm energy kilojoules, joules; chemist unit for heat: calorie kJ, J; cal pressure millimeters of mercury, Pascals, kilopascals, atmospheres mmHg, Pa, kPa, atm density temperature mass__ volume degree Celsius; g__; mL Kelvin o C; g__ cm3 K 1 km = 1 000 m 1 m = 100 cm 1 cm = 10 mm 1 m = 1 000 mm 1 kg = 1 000 g 1 g = 100 cg 1 cg = 10 mg 1 g = 1 000 mg 1 kL = 1 000 L 1 L = 1 000 mL 1 dm3 = 1 L 1 mL = 1 cm3 Dimension Analysis Dimensional Analysis (also called Factor-Label Method or the Unit Factor Method) is a problem-solving method that uses the fact that any number or expression can be multiplied by 1 one without changing its value. It is a useful technique. The only danger is that you may end up thinking that chemistry is simply a math problem - which it definitely is not. Unit factors may be made from any two terms that describe the same or equivalent "amounts" of what we are interested in. For example, we know that 1 inch = 2.54 centimeters Note: Unlike most English-Metric conversions, this one is exact. There are exactly 2.540000000... centimeters in 1 inch. We can make two unit factors from this information: Now, we can solve some problems. Set up each problem by writing down what you need to find with a question mark. Then set it equal to the information that you are given. The problem is solved by multiplying the given data and its units by the appropriate unit factors so that only the desired units are present at the end. (1) How many centimeters are in 6.00 inches? (2) Express 24.0 cm in inches. You can also string many unit factors together. (3) How many seconds are in 2.0 years? (4) Convert 50.0 mL to liters. (This is a very common conversion.) 2 (5) What is the density of mercury (13.6 g/cm3) in units of kg/m3? The mole It’s virtually impossible to conduct reactions one atom or molecule at a time, so the mole provides us with a convenient number of particles (atoms or molecules) to work with. One mole of anything is 6.02 x 1023 of those things, just like a dozen of something is 12 of those things. The molar mass of an element is the atomic mass of that element (on the periodic table) in grams. For example, a mole of carbon atoms is 6.02 x 1023 atoms and has a mass of 12.0 grams. A mole of water molecules is 6.02 x 1023 molecules and has a molar mass (or formula mass) of 18.0 grams (1.0 + 1.0 + 16.0) Dimensional analysis enables us to do conversions between moles, grams and particles. a) 2.5 moles of Mg are how many grams? Grams = 24.3 grams x 2.50 moles = 60.8 grams 1 mole b) 2.41 x 1024 atoms is how many moles? Moles = 1 mole x 2.41 x 1024 atoms = 4.00 moles 6.02 x 1023 atoms c) 120.3 grams of Ca is how many atoms? Atoms = 6.02 x 1023 atoms x 1 mole Ca x 120.3 grams = 1.81 x 1024 atoms 1 mole Ca 40.1 grams QUIZ: Question 1 Question 2 Question 3 How many millimeters are present in 20.0 inches? The volume of a wooden block is 6.30 in3. This is equivalent to how many cubic centimeters? A sample of calcium nitrate, Ca(NO3)2, with a formula mass of 164 g/mol, has 5.00 x 1027 atoms of oxygen. How many kilograms of Ca(NO3)2 are present? Answers: (1) 508 mm (2) 103 cm3 (3) 227 kg 3 Matter and Change Physical means of separating a mixture include: filtration, evaporation, using known freezing points and boiling points to separate different liquids, distillation (boiling off the liquid to leave the solid component, and then condensing the vapor back to the liquid state). Physical states of matter: Solid: particles packed very tightly together, particles are “fixed” in position relative to each other Liquid: particles still very close together but particles can move around each other Gas: particles very far apart from each other Physical changes involve changes in physical state (solid ↔ liquid ↔ gas: melting, boiling, condensation, freezing), cutting or crushing a large sample into smaller pieces, dissolving in an appropriate solvent. Chemical changes involve the rearrangement of the atoms of one or more substances to form one or more new substances. Physical properties of matter include: density, color, physical state at a given temperature, boiling point, freezing point, odor, malleability, brittleness, hardness, solubility in a given solvent, crystal shape. Chemical properties involve how the substance behaves in the presence of another substance: Does the substance give up electrons easily to another substance? Does the substance take electrons from another substance? When the substances are mixed together does a new substance form? Does a new substance form when the two substances are heated together? Does the substance combine readily with oxygen gas to form new substances? 4 Atomic Structure Atom: smallest particle of an element; composed of protons, neutrons, and electrons. Atomic Particle Location Electrical Charge Mass (atomic mass units) proton nucleus of atom +1 1 a.m.u. neutron nucleus of atom 0 1 a.m.u. electron outside of nucleus ─1 0 a.m.u. (mass is too small to be significant) In a neutral atom, the number of protons equals the number of electrons. Number of protons + number of neutrons = mass number of the atom (atomic mass in a.m.u.) The atomic number of an element is determined by the number of protons in every atom of that element. If “X” represents the symbol of the element, the mass number is written as a superscript at the upper left of the symbol, and the atomic number is written as a subscript at the lower left of the symbol: mass number X atomic number Isotopes: atoms of the same element that have the same atomic number (number of protons) but have different numbers of neutrons and therefore different mass numbers. and 157N are isotopes of nitrogen. Both have atomic number “7” but because the different numbers of neutrons, the atoms have different mass numbers, 14 and 15 respectively. 14 7N Another way of writing isotopes is to use the element symbol followed by the mass number: N-14, and N-15. Nuclear Chemistry Nuclear reactions are reactions that involve the nucleus of the atom. In all nuclear reactions, both mass and charge are conserved, meaning that the total mass and charge before a nuclear reaction will equal the total mass and charge after the reaction. 5 To the left is an example of a nuclear reaction, “fission,” where a large nucleus is split into 2 smaller nuclei. The numbers to the upper left of the symbol are mass numbers. The numbers to the lower left and the number of positive charges Nuclear reactions often involve the following, either as products or reactant: Alpha particle – a helium nucleus Beta particle – an electron Gamma Ray – a high energy wave Positron – exactly like an electron, but positively charged. Below are examples of radioactive decay, involving the above, where an unstable nuclide becomes another nuclide. Percent Composition There are two different ways to describe the composition of a compound: in terms of the number of its constituent atoms (like C2H6) and in terms of the percentages (by mass) of its elements. When showing the constituent atoms of a molecule, you can either show the chemical formula, which shows the real number of atoms in the molecule, like C2H6, or show the empirical formula, which merely shows their relative amounts in a substance, so the above molecular formula would be expressed as CH3. You can describe the composition of a compound in terms of the weights of its constituent elements by determining the percent composition of particular elements in the molecule. To calculate percent compositions, you would find the weight of each constituent atom, then figure out what percent of the total molecular weight it makes up. Consider ethanol, C2H5OH. Taking 6 subscripts into consideration, you have 2 moles of carbon, 6 moles of hydrogen (5 + 1), and 1 mol of O. Now convert moles into grams for each constituent element as well as for the entire molecule: Mass of C = 2 Mass of H = 6 Mass of O = 1 12.01 = 24.02 g 1.01 = 6.06 g 16.00 = 16.00 g Mass of 1 mol of C2H5OH = 46.08 g Now use the formula you learned above to find the percent compositions of the constituent elements: Mass percent of C: 100% = 52.14% Mass percent of H: 100% = 13.15% Mass percent of O: 100% = 34.77% Atomic Theory Dalton, Thompson and Rutherford and relevant experiments in determining atomic structure Dalton (1808)- Based on experimental results from previous century, devised scheme for each element being distinct, indivisible ("atomos" in Greek), and with a unique set of properties, including being very small, and combing with other atoms in distinct ways. He studied the ratios in which elements combine in chemical reactions. He discovered that all elements are composed of tiny invisible particles called atoms (an element is composed of only one kind of atoms, and a compound is composed of particles that are chemical combinations of different kinds of atoms); He came up with the first atomic theory which consisted of the following points Atoms of the same element are identical; The atoms of any one element are different from those of any other element; Atoms of different elements can physically mix together or can chemically combine to one another in simple whole number rations to form compounds; Law of Definite Proportions Atoms can combine in different ratios to form different compounds. Law of Multiple Proportions When reactions occur atoms rearrange, they are not created nor destroyed. Law of Conservation of Mass Thomson, Joseph John (J.J.)(1890's)- He performed experiments that involved passing electronic current through gases at low pressure. He knew that opposite charges attract and like 7 charges repel, so proposed that a cathode ray is a stream of tiny, negatively charged particles moving at high speeds.. Cathode Ray Tube Experiment Since these particles were smaller than atoms, but seemed to come from them, they must be subatomic parts. He concluded that electrons must be parts of atoms of all elements So, he discovered the first subatomic particle (electron), the atom is no longer indivisible, and developed the "plum pudding" (or "chocolate chip cookie") model. Plum Pudding Model of the Atom Rutherford, Ernest(1900's)- Discovered atoms are not homogeneous, very small part (the nucleus) is hard and heavy, electrons on the outside are very light and occupying the vast majority of the space of the atom. All from the gold foil experiment in which he directed a narrow beam of alpha particles at a very thin sheet of gold foil and noticed that the alpha particles should have passed easily through the gold, with only a slight deflection due to the positive charge thought to be spread out in the gold atoms. Gold Foil Experiment 8 He concluded that most of the alpha particles passing through pass through the gold foil because the atom is mostly empty space. The mass and positive charge is concentrated in a small region of the atom. Rutherford called this region the nucleus. Particles that approach the nucleus are greatly deflected. Solar System Model of the Atom Electrons in Atoms Soon after Rutherford disclosed his model of the atom it was discovered that electrons may “jump” from one orbit to a higher one. Later it was discovered that this jump happens ONLY IF the electron absorbs ENOUGH energy to make the jump. This is known as quantized energy. The Energy released when the electron fell back to the ground state was just as specific as the energy absorbed by the atom. In order to truly understand what is going on in the atom you must first understand electromagnetic radiation. Electro Magnetic Radiation Electromagnetic Radiation is energy in the form of waves. The electromagnetic spectrum is the range of all possible electromagnetic radiation. Also, the "electromagnetic spectrum" (usually just spectrum) of an object is the frequency range of electromagnetic radiation that it emits, reflects, or transmits. The electromagnetic spectrum, shown in the chart, extends from just below the frequencies used for modern radio (at the long-wavelength end) to gamma radiation (at the short-wavelength end), covering wavelengths from thousands of kilometers down to fractions of the size of an atom. Electromagnetic energy at a particular wavelength λ (in vacuum) has an associated frequency and photon energy E. Thus, the electromagnetic spectrum may be 9 expressed equally well in terms of any of these three quantities. They are related according to the equations: wave speed (c) = frequency x wavelength c = and E = h or where: c is the speed of light, 299792458 m/s h is Planck's constant, . . So, high-frequency electromagnetic waves have a short wavelength and high energy; lowfrequency waves have a long wavelength and low energy. Bohr Atom Niels Bohr quickly seized upon this used it to propose a quantized description of the atom. 1. Bohr proposed that while circling the nucleus of the atom, electrons could only occupy certain discrete orbits, that is to say energy levels. Bohr used Max Planck's equations describing quanta of radiation to determine what these discrete orbits would have to be. As long as electrons stay in these energy levels, they are stable. 2. Further, Bohr said electrons give or take energy only when they change their energy levels. If they move up, they take energy (say from light), and if they move down, they release energy. This energy itself is released in discrete packets called photons. 3. Furthermore, Bohr also said that an electron which is not in its native energy level (in other words, which has been excited to a higher energy level) always has to fall back to its original, stable level. Electrons may be excited by heat, light, electricity, or any other form of energy. The photons released by excited electrons returning to their normal energy levels accounts for the colors we see in flame tests, fireworks, any fire such as that in a fireplace or a lit match, and in the colors of our clothes (the electrons of the atoms in dye molecules are excited by light energy). Bohr Model of the Atom Bohr interpreted the emission lines in the spectra of gases as formed by the transitions of electrons from a higher to a lower energy level. Different elements have different allowed energy levels, and so released different energies during transitions. Release of photons in the visible range provide colors 10 unique to that element Hydrogen Spectrum Using the Bohr Atom Bohr also assumed that the electron can change from one allowed orbit to another The Periodic Table Facts about the Periodic table The periodic table was discovered in 1869 by Dimitri Mendeleev. The periodic table can be divided into groups and periods. The vertical columns are called GROUPS or “families” and are numbered left to right as 1-18. The horizontal rows are called PERIODS and are numbered top to bottom as 1-7 on the current periodic table. Elements in the same group or family have similar chemical and physical properties. Group 1 is the ALKALI METAL group. Group 2 is the ALKALINE EARTH METAL group. Groups 3-12 are the TRANSITION METAL groups. These metals often form compounds that are colored (red, green, yellow, etc.) 11 Group 13 is known as the BORON group. Group 14 is called the CARBON group. Group 15 is the NITROGEN group. Group 16 is the OXYGEN group. Group 17 is known as the HALOGEN group. Group 18 is the NOBLE GAS group. Electron Configuration The location of an electron within an atom is described by the 4 different quantum numbers, which represent the primary energy level (‘n’), the energy sublevel (described as s, p, d, or f orbital types), the orientation of the orbital (s sublevels have one orientation, p sublevels have 3, d sublevels have 5, and f have 7). The location of the atom is governed by three principles: Aufbau principle – an electron will occupy the lowest energy level available to it. Hund’s Rule - In p, d, and f sublevels, and electron will occupy and empty orbital if one is available, and will double up with another electron only if all orbitals are occupied. Pauli Exclusion Principle - Two electrons occupying the same orbital must have opposite spins. Orbital Notation The orbital notation for sodium: 1s Electron Configuration Notation for sodium: 2s 2p 2p 2p 3s 1s22s22p63s1 Noble Gas Notation for sodium: [Ne] 3s1 Questions: How many electrons can fit into the n=4 principle energy level? 12 Write the orbital notation, electron configuration notation, and Noble Gas notation for electrons in different elements Determine an element based on its electron configuration or Noble Gas notation Explain the 3 rules that govern electron configuration Chemistry Midterm Exam Review 2015 Unit 1. Matter and Change Name _____ Date _____/______/______ 1. Identify each of the following as an element, a mixture, or as a compound: a) iced tea b) ice c) table sugar d) silver 2. Classify each of the following as either a physical change or as a chemical change: a) bending a piece of glass b) melting an ice cube c) cooking a steak d) cutting grass e) burning wood f) sugar dissolving in water g) boiling water 3. Give the correct symbols for each of the following elements: a) sodium e) copper b) aluminum f) magnesium c) chlorine g) iron d) sulfur h) nitrogen 4. List the three common phases or states of matter: _________________________, ___________________________, and _________________________. 5. Which of the following is NOT a physical property of matter? _______________________ density texture color flammability odor malleability melting point luster boiling point 13 Unit 2. Measurement 1. Rank these measurements from the smallest to the largest: (hint: change all numbers to scientific notation) a) 5.3 x 104 m d) 0.005 7 m b) 7.7 x 103 m e) 5.1 x 10-3 m c) 4.9 x 10-2 m f) 0.072 m Correct order: 2. For each of the following pairs of units, which is the larger unit? a) centigram or milligram b) liter or centiliter c) calorie or kilocalorie d) millisecond or centisecond e) milliliter or kiloliter f) mm3 or m3 Show all work.. units on every number, cancel units where possible. 3. An object has a volume of 3.5 cm3 and a mass of 27.2 g. What is the density of the object? 4. What is the volume of an object having a density of 1.05 g/mL and a mass of 4.85 mL? Show all work. units on every number, cancel units where possible, pay attention to sig. figs and units in answer. 5. Write the following measurements in scientific notation: a) 572.5 km b) 0.005 725 m 6. Write the following measurements in standard notation: a) 4.45 x 10-3 g b) 4.45 x 107 mg 14 Unit 4. Atomic Structure and Nuclear Chemistry 1. Complete the following table – assume neutral atoms. Use your periodic table as a reference. Element K # protons Mass # # electrons Atomic # # neutrons 31 20 12 2. For each of the following atoms, give the atomic number and mass number: a) 36 18Ar atomic #: mass #: b) 79 35Br atomic #: mass #: Complete the following sentences: 3. The atomic number is determined by the number of 4. In a neutral atom, the number of in an atom. equals the number of 5. The mass number is determined by the total number of . and . 6. The difference in mass of isotopes of the same element is due to different numbers of ________________________ in the nucleus of the atoms. _____ 7. The major weakness in Rutherford’s model of the atom was that A) negatively charged electrons orbited the positively charged nucleus B) there was no explanation of why there were no neutrons in the nucleus C) there was no explanation of why the negatively charged electrons didn’t fall into the positively charged nucleus D) there was a dense positively charged nucleus Match the atomic model with the appropriate scientist or theory name _____ 8. John Dalton A) dense positively charged nucleus with negatively charged electrons moving around the nucleus _____ 9. Ernest Rutherford B) positively charged sphere in which negative charges are uniformly distributed 15 _____ 10. J.J. Thomson C) tiny, indivisible bits of matter; different kinds of matter made of different kinds of atoms _____ 11. Niels Bohr D) tiny invisible spherical bits of matter having a uniform density and which combine in definite proportions to form compounds _____ 12. ancient Greeks E) dense positively charged nucleus with negatively charged electrons moving at fixed distances from the nucleus; electrons could move from one level to a higher level if sufficient energy is absorbed _____ 13. modern atomic theory F) dense positively charged nucleus around which electrons move in areas of highest probability Unit 7 Periodic Trends 1. In the periodic table, a column of elements is known as a 2. The alkali metals are located in Group . 3. The alkaline earth metals are located in Group 4. The halogens are located in Group 5. The noble gases are located in Group . . . . 6. The element located in Group 15, Period 4 is . Two characteristics of metals are 7) 8) Two characteristics of nonmetals are 9) 10) 11. If an atom gains or loses one or more electrons, an 12. Atoms of an element with different numbers of neutrons are called _____of that element. 13. An atom with 17 protons, 20 neutrons, and 16 electrons has a mass number of ______, a charge of _______, and the name of the element is __________ 16 is formed. Given the elements Ga, Cl, and Al, put them in order of: 14. increasing atomic number 15. increasing atomic size 16. increasing first ionization energy 17. increasing electron affinity 18. increasing electronegativity Unit 6. Electrons in Atoms _____ 1. The light energy emitted by an “excited” electron is A) lost as the electron returns to a higher energy level B) lost as the electron returns to a lower energy level C) gained as the electron returns to a lower energy level D) gained as the electron returns to a higher energy level _____ 2. When an element is heated sufficiently to excite its electrons and the emitted light is passed through a diffraction grating or a prism, you would expect to observe A) a continuous spectrum C) a single color B) a bright-line spectrum D) white light _____ 3. A single burst of light is released from an atom. Which statement BEST explains what happened in the atom? A) An electron pulled a neutron out of the nucleus. B) An electron moved from a higher energy level to a lower energy level. C) An electron changed from a particle to a wave. D) An electron pulled a proton out of the nucleus. 17