STM124 - Midterm Reviewer 2.pdf

STM 124 (General one or more other substances. Chemistry 1) REVIEWER ➔ Ability to undergo chemical changes. FO...

STM 124 (General one or more other substances. Chemistry 1) REVIEWER ➔ Ability to undergo chemical changes. FOR MIDTERM ➔ Example: rusting of iron, EXAMINATION flammability, acidity and/or ph level. d. Physical Properties ➔ Does not involve chemical Unit 1 Lesson 1: Properties reactions. of Matter, Classification of ➔ Observed without changing Matter, & Simple the sample’s composition. ➔ Example: states of matter Separation Techniques (solid, liquid, gas), hardness, shape, texture, boiling and melting point. Chemistry ➔ Is the scientific study of the B. CLASSIFICATION OF MATTER properties and behavior of matter. a. PURE SUBSTANCES A. MATTER ➔ Has a fixed composition. ➔ Is anything that has mass and ➔ Properties do not change. occupies space. ➔ Cannot be broken to other pure ➔ Can change from one form to substances by any physical another. processes. ➔ Has three states: solid, liquid, and gas. 1. Elements ➔ Can be naturally occurring (from ➔ A substance that consists of nature) or synthetic (laboratory or identical atoms. factories). ➔ Can be found on the periodic table of elements. a. Chemical Change ➔ Example: hydrogen, oxygen, ➔ Conversion of a substance carbon, titanium. into another, hence, its chemical composition 2. Compounds changes. ➔ Composed of two or more ➔ Example: iron rusting, wood elements in a fixed ratio by burning, metabolism, or mass. cooking an egg. ➔ Characterized by its formula, which gives us the ratios of b. Physical Change the compound’s constituent ➔ Change without alterations in elements and identifies each the substance’s chemical element by its atomic composition. symbol. ➔ Usually a change in physical ➔ Example: water or H2O (two appearance/state. hydrogen atoms and one ➔ Example: crushing a can, oxygen atom), carbon melting, boiling, mixing dioxide or CO2 (one carbon salt/sand/sugar with water. atom and two oxygen atoms). c. Chemical Properties ➔ Ability of a substance to a. MIXTURES combine with or change into 1  ➔ A combination of two or more pure ➔ Example: desalination of substances. water ➔ The pure substances can be present in any mass ratio. c. Crystallization ➔ The formation of pure solid 1. HOMOGENEOUS MIXTURE particles of a substance from ➔ Has uniform composition. an impure mixture. ➔ Has a single phase. ➔ On adding a solid substance ➔ Example: saltwater, ink, in a liquid and stirring it, the honey, brass, air solid dissolves in the fluid. ➔ Saturation point: a point 2. HETEROGENEOUS where no solid dissolves in MIXTURE the liquid. ➔ Does not blend smoothly ➔ Example: crystallization of throughout. sugar from an aqueous ➔ Individual substances remain solution distinct. ➔ Example: freshly squeezed d. Sublimation juice, oil and water, smog, ➔ Solid changes to vapor pizza, bubble/milk tea, without passing through the blood). liquid phase. ➔ Can be used to separate two C. SIMPLE SEPARATION TECHNIQUES solids present in a mixture ➔ Used to separate the constituents of when one of the solids a mixture and are usually done to sublimes while the other remove unwanted materials or does not. obtain useful components. ➔ Example: dry ice, moth balls, air fresheners,. a. Filtration ➔ Uses a porous barrier to e. Paper Chromatography separate a solid from a ➔ Separates the components liquid. of a mixture based on the ➔ Based on particle size. ability of each component to ➔ Filtrate: the liquid that travel across a surface. passes through the filter ➔ Mobile phase: the material paper. that allows the substance to ➔ Residue: The solid trapped travel (e.g. alcohol). in the filter paper. ➔ Stationary phase: the ➔ Example: brewing coffee, material where the water filtration substance travels through (e.g. filter paper). b. Simple Distillation ➔ Example: separating ➔ Separates liquid from a pigments of colored ink. mixture of liquids by heating up the solution into gas and f. Evaporation condensing the wanted ➔ Used to separate a solution substance back into liquid of a solvent and a soluble form. solid. ➔ Based on differences in ➔ Separates solids from boiling points. liquids. ➔ Distillate: the liquid obtained ➔ Solute: the solid particle; from the condensation of what is left behind. vapors in distillation. ➔ Solvent: the liquid portion; what evaporates. 2  ➔ Example: Salt collected from saltwater mixture. 1. John Dalton ➔ Proposed the theory of the g. Mechanical Picking particle nature of matter. ➔ Separates solids from solids. ➔ He is regarded as the father ➔ Based on the difference in of atomic theory. physical characteristics like sizes and shapes of the a. Dalton’s Atomic Theory substances in a mixture. 1. All matter is composed of ➔ Example: removing husk tiny, indivisible particles particles from rice grains called atoms. 2. Each element has unique h. Magnetic Separation atoms, different from atoms ➔ Makes use of the magnetic of other elements. property of one component. 3. Atoms are neither created ➔ Ferromagnetic substances nor destroyed. such as iron, cobalt, and 4. Atoms of different elements nickel are separated from combine with each other in nonmagnetic substances certain whole number using magnetic separators. proportions to form ➔ Example: removing iron compounds. filings from sulfur powder b. Revisions of Dalton’s Atomic i. Decantation Theory ➔ Separation of liquid from 1. Atoms are not indestructible. They solid and other immiscible still consist of smaller particles. (non-mixing) liquids, by 2. The atoms of one element may differ removing the liquid layer at in mass. They are identical, the top from the layer of however, in some basic respects. solid or liquid below. ➔ The process can be carried c. Dalton’s Model out by tilting the mixture after ➔ His atomic theory is universally pouring out the top layer. accepted now as our current view of ➔ Example: Separation of oil matter. and water, removing water ➔ Laws following the atomic theory: from rice. 1. The law of conservation of mass states that there is no Unit 2 Lesson 2: The detectable change in mass during an ordinary chemical Development of the Atomic reaction. Model, The Periodic Table, 2. The law of constant composition states that a and Nomenclature chemical compound always contains the same elements A. DEVELOPMENT OF THE ATOMIC in the same proportion by MODEL mass. - Example: H2O is Leucippus and Democritus always H2O. If the ➔ Proposed that matter is composed composition of tiny particles that cannot be changes, it is not the divided. same compound ➔ They named it atomos, which anymore. means indivisible or uncuttable in 3. The law of multiple Greek. proportions states that if 3  two elements can be combined to form several 5. Quantum (Mechanical) Model possible compounds, then (Erwin Schrödinger) the ratios of the masses of ➔ Electrons move in waves around the the second element will be electron cloud. ratios of small whole numbers. B. SUBATOMIC PARTICLES - Example: If O has a 1. The electron is a negatively mass of 12g in CO, charged particle discovered by J.J. then in CO2, O has a Thomson in 1897. mass of 24g. 2. The proton is a positively charged particle discovered by Ernest 2. Plum Pudding Model (Joseph Rutherford in 1917. John Thomson) 3. The neutron is a subatomic particle ➔ An atom is a massive positively that bears no electric charge. It was charged blob embedded with discovered by James Chadwick in negatively charged electrons. 1932. Isotopic Notation by Henry Moseley ➔ Atomic Number (Z): number of protons. ➔ Atomic Mass (A): number of protons and neutrons. Figure B.1. Plum Pudding Model 3. Nuclear Model (Ernest Rutherford) ➔ An atom is made up of dense, positively charged nucleus surrounded by electrons. Figure B.4 - Atomic and Mass Number C. ATOMS, MOLECULES, IONS, AND ISOTOPES a. Atom ➔ The smallest particle of an element. b. Molecule Figure B.2. Nuclear Model ➔ Forms when two or more 4. Planetary Model / Bohr’s Model atoms combine. (Neils Bohr) c. Ions ➔ Electrons move around the nucleus ➔ Atoms that have electrical in orbits, like how planets move charges. around the sun. ➔ Some atoms can either gain or lose electrons; the number of protons never change in an atom. 1. Anions ➔ When an atom gains an electron or more, it becomes negatively charged. Figure B.3. Planetary Model 4  ➔ Applicable for - Elements are arranged nonmetals. according to atomic number. 2. Cations - Elements are organized into ➔ When an atom loses groups (have similar an electron or more, it physical/chemical becomes positively properties). charged. ➔ Applicable for metals. E. NOMENCLATURE: CHEMICAL 3. Monatomic ion NAMING AND FORMULA WRITING ➔ An ion composed of only one atom. Chemical formula 4. Polyatomic ion ➔ Shorthand symbol for ➔ An ion composed of compounds, showing the more than one atom. ratio of the atoms present in the compound. ➔ Empirical formula: shows the simplest ratio of atoms in the compound. ➔ Molecular formula: shows the actual ratio of atoms that comprise a molecule of that compound. ➔ Example: - Hydrogen Peroxide Molecular Formula: H2O2 or 2:2 Empirical Formula: Figure B.5 - Atoms, Molecules, and Ions HO or 1:1 - Glucose d. Isotopes Molecular Formula: ➔ Two atoms with the same C6H12O6 or 6:12:6 atomic number (number of Empirical Formula: protons) but different atomic CH2O or 1:2:1 mass due to their differences in the number of neutrons. I. Metal-Nonmetal Binary (Ionic) Compound Formula Consider charges. The Writing: charges become the subscripts of the other element using the criss cross method. - Zn2+ and O2- Since the elements have a charge of 2+ and 2-, they cancel Figure B.6 – The Three Isotopes of Hydrogen each other out. Hence, do not use subscripts. D. THE MODERN PERIODIC TABLE Formula: ZnO ➔ Dmitri Mendeleev is known to be Name: Zinc Oxide the father of the modern periodic - Na+ and S2- Using the criss cross table. method, charge of Na ➔ The periodic law: will be subscript of S, charge of S will be 5  subscript of Na. parentheses since it has two Formula: Na2S elements as one. Name: Sodium Sulfide Mg(ClO)2 - Magnesium Naming: Metals with Fixed Charges: Hypochlorite Metal + Nonmetal—ide Ex: Fe2+ and PO43- Ex: Zinc + Oxygen = Zinc Oxide Using the Criss Cross method, Iron now will have a subscript Transition Metals with two or of 3 while phosphate will have more charges: a subscript of 2, enclosed in parentheses. - IUPAC Name: Metal (Roman Numeral of Charge) Ion Fe3(PO4)2 - Iron (II) Phosphate or Ferrous Phosphate - Classical Name: - Classical name of Naming Name depends on the number metal—ous (lower of of oxygen atoms and NOT charge) Polyato- from the charges. - Classical name of mic metal—ic (higher Ions: According to the number of charge) oxygen atoms (smallest to largest) these are the rules in Ex: Copper can have a charge naming polyatomic ions: of 1+ and 2+. 1. Hypo_____ite Cu+ = Copper (I) Ion or 2. _____ite Cuprous Ion 3. _____ate 4. Per_____ate Cu2+ = Copper (II) Ion or Cupric Ion Ex: - ClO- = Hypochlorite Ex: Copper and Chlorine - ClO2- = Chlorite compound. Since Copper can - ClO3- = Chlorate have two charges, the formula - ClO4- = Perchlorate can be either of the two: Chemic Name of Metal (First Element) CuCl = Copper (I) Chloride or al + Name of Polyatomic Ions, Cuprous Chloride Naming following the rules given in the row above. CuCl2 = Copper (II) Chloride or Cupric Chloride Ex: KNO2 II. Compounds with Polyatomic - K+ and NO2- Ions - Potassium Nitrite Oxyanions - Anions composed of an III. Nonmetal-Nonmetal Binary element and oxygen. (Molecular) Compounds Formula Similar to Ionic Compounds, Formula Follow the number of atoms. Writing: however, if the oxyanion is Writing: given a subscript, enclose it in Ex: 2 atoms of Nitrogen and 5 a parenthesis and put the atoms of Oxygen = N2O5 subscript outside the parentheses. Naming: Use Greek Prefixes: 1 - Mono Ex: Mg2+ and ClO- 2 - Di 3 - Tri Use the Criss Cross Method 4 - Tetra and enclose CIO in 5 - Penta 6  6 - Hexa Acid 7 - Hepta 8 - Octa Name: If the anion ends in -ite, 9 - Nona it is changed to -ous followed 10 - Deca by the word “acid.” If the name 11 - Undeca ends in -ate, it is changed to 12 - Dodeca -ic followed by the word “acid.” For the first element, do not use mono if the number of V. Bases atoms is only 1. Simply just use the name of the element. Formula The metal will always have NO Writing: SUBSCRIPT since OH has a Ex: negative charge. - CO2 = Carbon Dioxide - N2O = Dinitrogen Enclose OH in parentheses Monoxide before writing down the - N2O5 = Dinitrogen subscript. Pentoxide Exception to the rules: Metal Nonmetal Formula Water: H2O Ammonia: NH3 K+ KOH Phosphine PH3 Al3+ OH+ Al(OH)3 Arsine: AsH3 Methane: CH4 Fe3+ Fe(OH)3 IV. Acids Naming: Metal Name + Hydroxide Binary Formula: H + Nonmetal Acids If Transition Metal, use IUPAC Formula Gaseous Aqueous or Classical Name + Hydroxide Form Form HF Hydrogen Hydrofluori Fluoride c Acid Formula Name HI Hydrogen Hydroiodic Iodide Acid KOH Potassium Hydroxide Name: - Gaseous Form: Al(OH)3 Aluminum Hydrogen + Hydroxide Nonmetal-ide Fe(OH)3 Iron (III) - Aqueous Form: Hydroxide Hydro-nonmetal-ic Acid Or Oxyacid Formula: H + Oxyanion s Ferric Hydroxide Formula Anion Chemical Name Name H2C2O4 Oxalate Oxalic Acid HlO Hypoiodite Hypoiodou Unit 3 Lesson 3: Chemical s Acid HlO2 Iodite Iodous Reactions and Equations Acid HlO3 Iodate Iodic Acid A. Chemical reaction HlO4 Periodate Periodic ➔ A process wherein at least one substance is produced 7  as a result of a chemical (aq) Aqueous Solution ( substance is in change. a solution with water) ➔ Chemical reactions follow the Law of Conservation of ∆ Heating process Mass, which states that atoms are neither created D. Writing Chemical Equations nor destroyed during a 1. Follow the rules in writing chemical reaction; they are chemical formulas of a just rearranged compound. 2. Some elements are diatomic B. Chemical equations molecules (H2, N2, O2, F2, Cl2, ➔ Represents chemical Br2, and I2) while others are reactions through symbolic monatomic gasses (He, Ne, Ar, representations. Kr, Xe). ➔ Uses formulas and symbols 3. If ∆ is found on top of the arrow, to describe the changes that it means that heat is required for have occurred in the the reaction to take place. reaction. 4. Law of conservation of mass: - Mass is neither created C. Features of a Chemical Equation nor destroyed. a. Reactants - substances to - Number of atoms in the the LEFT of the arrow. reactants must be b. Products - substances to EQUAL to the number of the RIGHT of the arrow; what atoms in the products. is formed during the reaction. Example: Sodium metal reacts with c. Subscripts - Number of chlorine gas to produce solid atoms of an element present sodium chloride. in a substance; FIXED, - Na(s) + Cl2(g) → NaCl cannot be changed. d. Coefficients - Number of E. Balancing Chemical Equations molecules of each ➔ Chemical equations need to substance. be balanced for us to be able to determine the amount of The following symbols are used in writing reactants required to chemical equations: produce a specific amount of product. → Separates reactants from the ➔ The number of atoms of products. each element in the reactants should be equal to “Yields,” “Produces,” “Forms,” or the number of atoms of each “Liberates.” element in the products. ⇌ Indicates that the reaction is ➔ If there are 2 hydrogen reversible. atoms in the reactants, then there should also be 2 + Separates reactants from each hydrogen atoms in the other. products. ↑ Forms a gaseous product. ➔ Use coefficients to balance equations. ↓ Forms a precipitate (liquid to solid). Example 1: Calcium reacts with (s) Solid hydrochloric acid to form calcium chloride (g) Gaseous (l) Liquid and hydrogen gas. 8  1. Write down the equation: of oxygen for both the products and the reactants. Ca + HCl(l) → CaCl2 + H2 - Divide 6 by 2 to get the coefficient for the reactant. 2. Count the number of atoms of each 6/2 = 3 element in both the reactants and - Divide 6 by 3 to get the the products. coefficient for the product. 6/3 = 2 Element Atoms in Atoms in Reactant Product Add the coefficients: Ca 1 1 Fe + 3O2 → 2Fe2O3 Note that for the products, Fe2O3 is H 1 2 considered as one. Hence, the Cl 1 2 coefficient should be placed before Fe and not in between the two 3. Do a trial and error process of elements. balancing out the atoms. Now, check if all elements are Note: balanced. In this case, Fe is still - Based on Step 1, balance unbalanced. The product now has 4 out Cl first since in the Fe atoms (multiply the coefficient to products, it is binded with the subscript), while the reactant Ca and in the reactants, it is only has 1 Fe atom. Simply add 4 as one with H. the coefficient of Fe in the reactants. - Since there are 2 Cl in the 4Fe + 3O2 → 2Fe2O3 products, add a coefficient Counting the number of atoms of of 2 before HCl in the each element, the equation is now reactants so that Cl will be balanced. balanced. Ca + 2HCl(l) → CaCl2 + H2 F. EVIDENCES OF A CHEMICAL - Now, check if there are any CHANGE more atoms that are 1. Change in color. unbalanced. 2. Change in temperature - If none, that is your final (heat is released or answer. If there is, continue absorbed). adding/changing the a. Endothermic coefficients. Reaction - heat is absorbed by the Example 2: Iron reacts with oxygen gas to substances (e.g., produce iron (III) oxide. photosynthesis and dissolving salt in 1. Fe + O2 → Fe2O3 water). 2. b. Exothermic Element Atoms in Atoms in Reaction - heat is Reactant Product released to the Fe 1 2 environment. (e.g., rusting, freezing of O 2 3 water, and nuclear fission). 3. Balance O first: 3. Evolution of gas (bubble - Since the atoms are 2 and 3 for formation. Oxygen, their least common 4. Formation of a solid multiple is 6. Thus, 6 should be (precipitate). the balanced number of atoms 5. Production of light. 9  - Types of Chemical Reactions 1. Combustion Reaction - Substance reacts with oxygen, releasing energy. - Products are always CO2 and H2O. Example: CH4(g)+ 2O2 (g) → CO2 (g) + 2 H2O (l) 2. Combination Reaction - Also called synthesis reaction. - Two or more substances form a single new substance. - A + B → AB Example: NO + O2 → NO3 3. Decomposition Reaction - A compound breaks down into two or more substances. - The “reverse of Combination Reaction.” - AB → A + B Example: CaCO3 → CaO + CO2 4. Single Displacement Reaction - A more reactive element replaces a less reactive element. - AB + C → AC + B (B was replaced by C) Example: Zn + CuSO4 → ZnSO4 + Cu 5. Double Displacement Reaction - There is an exchange of positive ions between two compounds. - AB + CD → AC + BD Example: 2 KI(aq) + Pb(NO3)2 (aq) → 2 KNO3 (aq) + PbI2 (s) 10 Practice Questions 4. Cobalt-60 is used medically for radiation therapy as implants and as an external See page 12 for the answer key. Good luck! source of radiation exposure. What is the atomic number, atomic mass, number of p+ 1. Fill in the blanks: , e- , n0 of this isotope? Ions Chemical Formula Chemical Name Ag+ Br- a. Atomic Number = 60, Atomic Mass = 27, p+ = 33, e- = 33 , n0 = 27 Fe2O3 b. Atomic Number = 27, Atomic Mass K+ PO43- = 27, p+ = 27, e- = 60 , n0 = 33 OH- Lithium Hydroxide c. Atomic Number = 33, Atomic Mass = 60, p+ = 60, e- = 60 , n0 = 27 Carbonic Acid d. Atomic Number = 27, Atomic Mass (NH4)2SO3 = 60, p+ = 27, e- = 27 , n0 = 33 Ba2+ S2- 5. Thallium-204 is used in survey meters by SnO2 schools, the military and emergency N2O5 management authorities. What is the P4S10 atomic number, atomic mass, number of p+ , e- , n0 of this isotope? Sulfur Hexafluoride No Ions NO2 a. Atomic Number = 123, Atomic Mass Nitrogen Trifluoride = 204, p+ = 123, e- = 123, n0 = 81 b. Atomic Number = 81, Atomic Mass N2O7 = 204, p+ = 81, e- = 81 , n0 = 123 Carbon c. Atomic Number = 81, Atomic Mass Tetrachloride = 204, p+ = 204, e- = 204 , n0 = 123 d. Atomic Number = 204, Atomic Mass 2. Write the balanced chemical = 81, p+ = 123, e- = 123 , n0 = 81 equation of the following: a. Hydrochloric Acid reacts with an aqueous solution of Barium Hydroxide to produce Water and Barium Chloride. b. When heated, aqueous sodium iodide reacts with liquid bromine to form aqueous sodium bromide and solid iodine. c. Calcium reacts with hydrochloric acid to form calcium chloride and hydrogen gas. 3. Balance the given chemical equations and determine the type of reaction that took place: a. C3H8 + O2 → CO2 + H2O b. Al + HCl → AlCl3 + H2 c. Na3PO4 + MgCl2 → NaCl + Mg3(PO4)2 d. C2H6 + O2 → CO2 + H2O e. H2 + O2 → H2O 11 Answers 4. d. Atomic Number = 27, Atomic Mass = 60, p+ = 27, e- = 27 , n0 = 33 1. Fill in the blanks: 5. b. Atomic Number = 81, Atomic Mass = Ions Chemical Formula Chemical Name 204, p+ = 81, e- = 81 , n0 = 123 Ag+ Br- AgBr Silver Bromide _______________________________________ Iron (III) Oxide / Fe3+ O2- Fe2O3 Ferric Oxide K+ PO43- K3PO4 Potassium Prepared by: Phosphate Andrea Jolie Chua Li+ OH- LiOH Lithium Hydroxide H+ CO32- H2CO3 Carbonic Acid Jhullia Francheska Ebol NH4+ SO32- (NH4)2SO3 Ammonium Sulfite Ba2+ S2- BaS Barium Sulfide Tin (IV) Oxide / Sn4+ O2- SnO2 Stannic Oxide Dinitrogen N2O5 Pentoxide Tetraphosphorus P4S10 Decasulfide SF6 Sulfur Hexafluoride No Ions NO2 Nitrogen Dioxide NF3 Nitrogen Trifluoride Dinitrogen N2O7 Heptoxide Carbon CCl4 Tetrachloride 2. Writing balanced chemical equations: a. 2HCl + Ba(OH)2 → 2H2O + BaCl2 b. 2NaI + Br2 → 2NaBr + I2 c. Ca + 2HCl → CaCl2 + H2 3. Balancing chemical equations and the types of reactions: a. C3H8 + 5O2 → 3CO2 + 4H2O - Combustion reaction b. 2Al + 6HCl → 2AlCl3 + 3H2 - Single displacement reaction c. 2Na3PO4 + 3MgCl2 → 6NaCl + Mg3(PO4)2 - Double displacement reaction d. 2C2H6 + 7O2 → 4CO2 + 6H2O - Combustion reaction e. 2H2 + O2 → 2H2O - Combination reaction 12