What You Need To Know for the Chemistry Regents ExamMr. Dunleavy Port Jervis High School misterd@hvc.rr.com What is the point of this packet? This review packet was assembled from NY State’s Core Curriculum, which outlines the material to be tested on the Regents exam. This is by no means a thorough review of the entire course. It is designed to be used with review sheets, past Regents exams and your Reference tables to help you prepare for the coming test. Emphasis is placed on key ideas that are stressed by the Core Curriculum. Additional space has been left for you to add your own notes. You cannot passively prepare for the Chemistry Regents. There are no shortcuts. You have to study, ask questions, analyze problems and come to review sessions to be thoroughly prepared. Be Prepared Students taking the exam should come with a 4-function or scientific calculator (not a graphing calculator), pen, and pencil. Reference Tables will be provided. Students are required to stay in the examination room for a minimum of 2 hours from the time the test is distributed. Exam Date:_________________________________
The Test This analysis is based on the last 6 exams (June 2004 – January 2006). It is possible the new exams will have minor modifications. The Chemisty Regents Exam is broken down into three sections: Part A: 35 mulitple choice questions from all units covered over the
Part C: 20 short answer questions, most broken down into smaller
Each question is worth 1 point, for a total of 85 points. The grading scale changes with each test, with a raw score of 48-51 translating to a passing grade of 65. The actual scoring scale varies with each exam, and there is no way to know what curve will be used before the exam begins. The June 2004 scale is included on the last page for your reference. Topics There are 12 specific chemistry topics covered on the test, as well as a 13th topic on general lab skills: The Atom (p3) Moles and Stoichiometry (p15) Nuclear Chemistry (p7) Solutions (p16) Bonding (p8) Kinetics and Equilibrium (p17) Matter (p10) Acids, Bases and Salts (p18) Energy (p12) Oxidation-Reduction (p19) The Periodic Table (p13) Organic Chemistry (p20)
1. The modern model of the atom has evolved over a long period of time through the work of many scientists.
2. Each atom is made of a positively charged nucleus with one or more orbiting, negatively charged electrons. 3. Protons and neutrons are found in the nucleus. 4. Protons have a positive charge, neutrons no charge, and electrons a negative charge. 5. The number of protons in an atom equals the number of electrons.
6. The mass of a proton is 1 amu. The mass of a neutron is 1 amu. The mass of an electron is almost 0 amu.
and neutrons.
7. Each electron in an atom has its own distinct amount of energy.
8. When the electron gains a specific amount of energy, it moves to a higher orbital and is in the "excited state". 9. When an electron returns from a higher energy state to a lower energy state, it emits a specific amount of energy usually in the form of light. This can be used to identify an element (bright line spectrum).
spectroscope.
10. The outermost electrons are called valence electrons. These affect the chemical properties of the element.
11. Atoms of the same element all contain the same number of protons.
12. Isotopes are atoms with equal numbers of protons but different numbers of neutrons.
13. The average atomic mass of an element is the weighted average of its naturally occurring isotopes.
14. When an atom gains an electron, it becomes a negative ion and its radius increases. 15. When an atom loses an electron, it becomes a positive ion and its radius decreases. 16. Electron-dot diagrams (Lewis structures) represent the valence electron arrangement in elements, compounds and ions.
17. Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. These values are based on an arbitrary scale.
Topic Two: Nuclear Chemistry 1. The stability of an isotope depends on the ratio of protons to neutrons in the nucleus.
2. Each isotope has a specific mode and rate of decay. (see Table N)
3. A change in the nucleus of an atom changes it to a new type of atom (i.e. a new element). This is called transmutation. Transmutation can occur naturally or artificially.
4. Spontaneous decay can involve the release of different particles from the nucleus.
5. Nuclear reactions include natural and artificial decay, nuclear fission and nuclear fusion.
6. Nuclear reactions can be represented by equations that include symbols which represent atomic nuclei (with mass number and atomic number), subatomic particles (with mass and charge) and emitted particles. 7. Energy from nuclear reactions comes from the very small fraction of mass that is lost – the reaction converts matter into energy.
8. The energy released from nuclear reactions is much greater than that released from chemical reactions. 9. The risks associated with using radioactive isotopes include biological exposure (which may cause radiation poisoning and cancer), long-term storage and disposal, and nuclear accidents. 10. Radioactive isotopes may be used in medicine (tracing chemical and biological processes), radioactive dating, industrial measurement, nuclear power, and detection and treatment of disease. Topic Three: Bonding 1. Chemical compounds are formed when atoms are bonded together.
2. Two major categories of compounds are ionic and molecular (covalent) compounds. 3. Compounds can be differentiated by their chemical and physical properties.
4. Atoms gain a stable electron configuration by bonding with other atoms.
5. Chemical bonds are formed when valence electrons are:
6. In multiple (double or triple) covalent bonds more than 1 pair of electrons are shared between two atoms.
7. Polarity of a molecule can be determined by its shape and the distribution of the charge.
8. The electronegativity difference between two bonded atoms can determine the type of bond and its polarity. 0.0 - 0.4 = non-polar covalent 0.4-1.7 = polar covalent 1.7+ = ionic 9. Bonding guidelines:
10. Intermolecular forces allow different particles to be attracted to each other to form solids and liquids.
11. Physical properties of a substance can be explained in terms of chemical bonds and intermolecular forces. These include conductivity, malleability, solubility, ductility, hardness, melting point and boiling point.
Topic Four: Matter, Phases and Gas Laws 1. Matter is classified as a pure substance or a mixture of substances.
2. A mixture is composed of two or more different substances that may be physically separated.
2. An element is a substance composed of atoms with the same atomic number. They cannot be broken down by chemical change. 3. A compound is two or more elements bonded together. It can only be broken down by chemical changes.
4. A physical change is one that results in the rearrangement of existing particles in a substance (ex: freezing, boiling). A chemical change results in the formation of different substances with different properties.
5. The three phases of matter are solid, liquid and gas. Each has its own properties.
6. A heating curve (or cooling curve) traces the changes in temperature of a substance as it changes from solid to liquid to gas (or gas to liquid to solid).
7. Heat of fusion (Hf) is the energy needed to convert one gram of a substance from solid to liquid. 8. Heat of vaporization (Hv) is the energy needed to convert one gram of a substance from liquid to gas. 9. Specific heat (C) is the energy required to raise one gram of a substance 1 degree (Celcius or Kelvin).
10. The combined gas law states the relationship between pressure, temperature and volume in a sample of gas.
11. An ideal gas model is used to explain the behavior of gases. A real gas is most like an ideal gas when it is at high temperature and low pressure.
12. The Kinetic Molecular Theory (KMT) for an ideal gas states that all gas particles:
13. Equal volumes of gases at the same temp and pressure have an equal number of particles. Topic Five: Energy 1. Energy can exist in different forms – chemical, electrical, electromagnetic, thermal, mechanical, nuclear.
2. The Law of Conservation of Energy states that energy can not be lost or destroyed, only changed from one form to another. 3. Heat is a transfer of energy (often but not always thermal energy) from a body of higher temperature to a body of lower temperature. 4. Temperature is a measure of the average kinetic energy of the particles in a sample. Temperature is NOT a form of energy and should not be confused with heat. 5. The concepts of kinetic and potential energy can be used to explain physical processes such as fusion (melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, and deposition. 6. Processes that are exothermic give off heat energy. This typically causes the surrounding environment to become warmer. 7. Processes that are endothermic absorb energy. This typically causes the surrounding environment to become colder.
Topic Six: The Periodic Table 1. The placement of an element on the Periodic Table gives an indication of the chemical and physical properties of that element. 2. Elements are arranged in order of increasing atomic number. 3. The number of protons in an atom (atomic number) identifies the element.
4. The atomic mass is the sum of protons and neutrons in the nucleus.
5. Isotopes of an element are identified by the sum of protons and neutrons.
6. Elements can be classified by their properties and their location on the Periodic Table as metals, non-metals, metalloids, and noble gases. 7. Elements may be differentiated by their physical properties.
8. Elements may be differentiated by their chemical properties.
9. Elements are arranged into periods and groups. 10. Elements of the same period have the same number of occupied energy levels.
11. Elements of the same group have the same valence configuration and similar chemical properties.
12. The succession of elements within a group demonstrates characteristic trends in properties. As you progress down a group:
13. The succession of elements within a period demonstrates characteristic trends in properties. As you progress across a group from left to right:
14. Some elements may exist in two or more forms in the same phase. These forms differ in their molecular or crystal structure, hence their different properties.
Topic Seven: Moles and Stoichiometry 1. A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can only be broken down by chemical means. 2. Chemical compounds can be represented by a specific formula and assigned a name based on the IUPAC system. 3. Types of chemical formulas include empirical, molecular, and structural.
4. All chemical reactions show a conservation of mass, energy and charge. 5. A balanced chemical equation represents conservation of atoms. 6. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. 7. The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula mass) equals the mass of one mole of that substance. 8. The percent composition by mass of each element in a compound can be calculated mathematically. 9. Types of chemical reactions include synthesis, decomposition single replacement, and double replacement. Topic Eight: Solutions 1. A solution is a homogeneous mixture of a solute dissolved in a solvent.
2. Ionic substances dissolve in polar solvents. The positive ion is attracted to the negative end of the polar molecule, as the negative ion is attracted to its positive end. 3. Concentration of a solution can be expressed as molarity (M), percent by volume, percent by mass, or parts per million (ppm). 4. Adding a solute to a solvent causes the boiling point of the solvent to increase and the freezing point to decrease. 5. A saturated solution exists in equilibrium – the rate of crystallization equals the rate of dissolving.
Topic Nine: Kinetics and Equilibrium 1. Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and orientation. 2. The rate of a chemical reaction depends on temperature, concentration, nature of the reactants, surface area and the presence of a catalyst. 3. Energy absorbed or released by a chemical reaction can be represented by a potential energy diagram. 4. The amount of energy released or absorbed during a chemical reaction is the heat of reaction.
5. A catalyst provides an alternative pathway for a chemical reaction. The catalyzed reaction requires a lower activation energy than the uncatalyzed reaction.
6. Entropy is a measure of the randomness or disorder in a system. A system with greater disorder has greater entropy. 7. Systems in nature tend to undergo changes towards lower energy and higher entropy. 8. Exothermic reactions that result in increased entropy are spontaneous. 10. At equilibrium the rate of the forward reaction equals the rate of the reverse reaction. 11. The measurable quantities of reactants and products remain constant at equilibrium. 12. LeChatelier’s principle can be used to predict the effect of stress on a system in equilibrium.
Topic Ten: Acids, Bases and Salts 1. Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius acids and bases are electrolytes. 2. An electrolyte is a substance which, when dissolved in water, forms a solution capable of conducting electricity. The ability to conduct electricity depends on the concentration of ions. 3. Arrhenius acids yield H+(aq) ions as the only positive ion in solution.
4. Arrhenius bases yield OH-(aq) ions as the only negative ion in solution.
5. In neutralization reactions an Arrhenius acid and an Arrhenius base react to form salt and water.
6. Titration is a lab process in which a volume of a solution of known concentration is used to determine the concentration of another solution. Titration is a practical application of a neutralization reaction. 7. There are other acid-base theories besides the Arrhenius theory. The Bronsted-Lowry theory states an acid is a H+ donor and a base a H+ acceptor. 8. The acidity or alkalinity of a solution can be measured by pH.
9. On the pH scale, each decrease of one pH unit represents a tenfold increase in H+ ion concentration.
Topic Eleven: Oxidation-Reduction (Redox) 1. An oxidation-reduction (redox) reaction involves the transfer of electrons (e-). 2. Reduction is the gain of electrons and decrease of oxidation number.
3. Oxidation is the loss of electrons and increase of oxidation number.
4. In redox the number of electrons lost is equal to the number of electrons gained. 5. Oxidation numbers/states can be assigned to atoms and ions. Changes in oxidation numbers indicate that a redox reaction has occurred.
6. An electrochemical cell can be either voltaic or electrolytic. 7. In an electrochemical cell oxidation occurs at the anode and reduction at the cathode.
8. A voltaic cell spontaneously converts chemical energy to electrical energy. 9. An electrolytic cell requires energy to produce a chemical change. This is called electrolysis.
Topic Twelve: Organic Chemistry 1. Organic compounds consist of carbon atoms which bond to each other in chains, rings and networks to form a variety of structures. 2. Organic compounds can be named with the IUPAC system. 3. Hydrocarbons are compounds that contain only carbon and hydrogen.
4. Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, amines, amides, and amino acids are categories of organic molecules that differ in their structures. 5. Functional groups give organic molecules distinct physical and chemical properties. 6. Isomers of organic compounds have the same molecular formula but different structures and properties. 7. In a multiple covalent bond, more than one pair of electrons are shared between two atoms. Unsaturated organic compounds contain at least one double or triple bond. 8. Types of organic reactions include: addition, substitution, polymerization, esterification, fermentation, saponification, and combustion. Topic Thirteen: Lab skills 1. Any standard chemistry lab procedure is fair game for the Regents. Specific skills that may be tested include:
Regents Examination in Physical Setting / Chemistry June 2004
Chart for Converting Total Test Raw Scores to Final Examination Scores (Scaled Scores)
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