Unit A: Matter and Energy in Chemical Change Notes PDF

Unit A: Matter and Energy in Chemical Change  Chapter 1: Understanding Matter Properties of Matter Physical properties: describe physical appearance and composition of a substance; can be observed without changing the chemical identity of the substance. Boiling point: temperature at which...

Unit A: Matter and Energy in Chemical Change  Chapter 1: Understanding Matter Properties of Matter Physical properties: describe physical appearance and composition of a substance; can be observed without changing the chemical identity of the substance. Boiling point: temperature at which substance changes from a liquid to a gas. Melting point: temperature at which a substance changes from a solid to a liquid. Malleability: ability to be rolled into sheets without crumbling.  Ductility: ability to be stretched into a wire without breaking. Colour: wavelength of light reflected. State: solid, liquid, or gas. Solubility: Ability to dissolve. Crystal formation: crystalline appearance. Conductivity: ability to conduct electricity. Magnetism: magnetic attraction between objects. Chemical properties: describe the reactivity of a substance; can only be observed through chemical reactions. Ability to burn: combustible. Flash point: temperature required to ignite. Behaviour in air: tendency to degrade, react, or tarnish.  Reaction with water: tendency to corrode or form new substances. Reaction with acids: corrosion, bubble formation, neutralization. Reaction to heat: tendency to react due to temperature changes. Reaction with litmus: indicates if the substance is acidic, basic, or neutral. Qualitative observations: observed using the senses - e.g. state, colour, texture, and shape. Quantitative observations: can be measured using the correct instruments – e.g. melting point, boiling point, solubility, and conductivity. Classification of Matter Food Chemistry Cooking can result in either a physical change (change in state but the chemical components remain the same) or a chemical change (resulting in the formation of different substances) Important in development of food preservation techniques such as drying, heating, freezing, fermentation, and chemical preservation Metallurgy Science of producing and using metals Until 3000 B.C., the only known metals were gold, copper, lead, silver, and iron Different properties of metals resulted in their uses (eg. gold was soft – used to make jewelry, copper was hard – used to make weapons and tools) Different processes were developed to improve the quality of metals such as annealing, smelting and the making of alloys. Alchemy Combination of science and magic Attempt to transform cheap metals into gold Contributed to development of chemistry by discovering elements and their properties (eg. Mercury), originating methods and procedures, and developing laboratory equipment (eg. glassware, distillation apparatus) Atomic Models Video - Atomic Models History 1. Aristotle belief that matter was composed of mixtures of fire, air, earth, and water (believed at the time to be the smallest piece) 2. Democritus matter made up of tiny “indivisible” particles (ie. could not be divided into smaller pieces) – Atoms 3. Robert Boyle measured relationships between volume and pressure of gases and concluded that gases were made up of tiny particles that group together to make different substances 4. Antoine Lavoisier measured masses of reacted and produced substances - discovered that mass is neither produced nor lost – Law of conservation of mass Atomic Models Video - Atomic Models History 5. John Dalton (1800 AD) – Billiard Ball Model All matter is made of small, indivisible particles called atoms. All atoms of an element have identical properties and atoms of different elements have different properties. Atoms of different elements combine in specific fixed ratios to form new substances. 6. J.J. Thomson (1900 AD) – Plum Pudding Model Produced a beam of particles in a vacuum tube Determined that this beam was made up of negatively charged particles – discovery of electrons Suggested the atom was a positively charged sphere with negatively charged electrons embedded in it. 7. Ernest Rutherford (1920 AD) Observed that when positively charged particles are directed at a piece of gold foil, most passed through, but a 1/10000 bounced back. Concluded that the atom mostly empty space with a small, dense, positively charged core called the nucleus, with electrons orbiting the nucleus. 7.5 John Chadwich The nucleus contains protons (p+) and neutrons (n0). 8. Neils Bohr (1930) Electrons orbit the nucleus at specific energy levels. When electrons of a given element fall from higher energy levels to lower energy levels, they release a particular colour of light. 9. The Quantum Mechanical Model of the Atom – Current Model of the Atom (1940 AD) The nucleus is composed of nucleons – protons and neutrons. Electron orbitals describe the probable locations of electrons. Each electron is a probability cloud of negative charge that surrounds the nucleus at a particular energy level. Video – TED Ed – How small is an atom? Chapter 2: Classification and Naming of Compounds Atomic Theory Atomic Theory An atom is the smallest part of an element that still has the properties of the element; composed of three types of subatomic particles: The protons (p+) and neutrons (n0) in the nucleus comprise 99.9% of an atom’s mass. Electrons surround the nucleus and comprise 99.9% of the volume of an atom.  atomic number = number of protons number of electrons = number of protons number of neutrons = rounded mass number - number of protons Electrons in energy levels closest to the nucleus have the least energy and are more tightly held in the atom. The first energy level holds 2 electrons, the second and third energy levels hold 8 electrons each, and the fourth energy level holds 18 electrons. The electrons in the outer energy level are called valence electrons. Innermost energy levels must be completely filled before electrons can be placed in the next energy level.  Energy level diagrams are used to represent atomic structure. Example Problem 1: Draw an energy level diagram for a sulfur atom. Example Problem 2: Draw an energy level diagram for a hydrogen atom. The Periodic Table Periodic Table Song The periodic table organizes all of the elements according to their chemical properties. Elements in the same period have the same number of occupied energy levels. Elements in the same group/family have the same number of valence electrons, thus similar chemical and physical properties. oAlkali metals (group 1): soft, shiny, silver in colour, highly soluble and reactive in water. oAlkaline-earth metals (group 2): shiny and silver, but not as soft or soluble as group 1. oHalogens (group 17): poisonous and highly reactive. oNoble gases (group 18): very unreactive due to complete valance shells. Number of valence electrons and occupied energy levels for main group elements: The elements on the periodic table can be categorized as metals, non-metals, or metalloids. Metals are silver or grey in colour, shiny, malleable, ductile, good conductors of heat and electricity, and most are solid at room temperature; found to the left of the staircase. Non-metals may be solid, liquid, or gas, and are poor conductors of heat and electricity; found to the right of the staircase. Metalloids have properties that are intermediate between metals and non-metals. Ions and Isotopes Isotopes: atoms of the same element that contain different numbers of neutrons, and therefor have a different mass number – e.g. hydrogen-2. The atomic molar mass on the periodic table is an average of all of the element’s isotopes. The most common isotope of each element is found by rounding the mass number found on the periodic table. Example problem 3: write the symbol for oxygen-18. How many neutrons are in an oxygen-18 atom? Ionization: the process by which an atom gains or loses electrons to form an ion. Cation: positively charged ion that forms when an atom (usually a metal) loses one or more electrons. E.g. a sodium ion loses an electron giving it a 1+ charge – now called a sodium ion and written as Na+. If the ion has more than one possible charge, indicate the charge in roman numerals after the name – e.g. Ni2+ is called a nickel(II) ion. Anion: negatively charged ion that forms when an atom (usually a non-metal) gains one or more electrons. E.g. a chlorine atom gains one electron giving it a 1- charge – now called an chloride ion and written as Cl-. Example Problem 4: write the chemical symbol and draw an energy level diagram for an oxide ion Example Problem 5: write the chemical name for N3- Example Problem 6: write the chemical name for Zn2+ Example Problem 7: write the chemical symbol and draw an energy level diagram for a lithium ion Example Problem 8: write the chemical name for Fe3+ Ionic Compounds The International Union of Pure and Applied Chemistry (IUPAC) is the body responsible for naming compounds. Ionic compounds form when electrons transfer from one atom to another so that each atom has a complete valence energy level/stable octet. The cation and anion are attracted to one another by charge, forming a neutral ionic compound joined by an ionic bond.  E.g. magnesium and oxygen combine to form magnesium oxide – this represents one formula unit. Rules for naming binary ionic compounds: 1. Name the cation first by using the element’s name (usually a metal). 2. Name the anion second by using the first part of the element’s name (usually a non-metal) and changing the ending to –ide. Example Problem 9: CaCl2(s) Example Problem 10: MgO(s) Example Problem 11: K2S(s) Rules for writing formulas for binary ionic compounds: 1. Write the symbol for the cation followed by the symbol for the anion. 2. Determine the total number of each ion required to balance the charges. 3. Use subscripts to indicate the ratio of cations to anions – no subscript is needed for a single atom. Example Problem 12: calcium bromide Example Problem 13: silver sulfide Example Problem 14: magnesium phosphide Example Problem 12: calcium bromide Example Problem 13: silver sulfide Example Problem 14: magnesium phosphide Multivalent elements: have more than one stable ion charge; the first charge listed on the periodic table is the most common one. Roman numerals in brackets after the element name indicate the charge of the ion – e.g. iron(II) bromide. When given only the chemical formula, use the anion charge to find the cation charge – e.g. FeBr2 is iron(II) bromide Example Problem 15: copper(II) chloride Example Problem 16: Fe2O3 Polyatomic ions: made up of several non-metallic atoms joined together by covalent bonds. Naming rules are the same as binary ionic compounds, except the ending of the polyatomic ion name is NOT changed. Brackets are used around polyatomic ions when a subscript is needed. Example Problem 17: Ca(NO2)2 Example Problem 18: ammonium carbonate Properties of ionic compounds include: High melting point/solid at room temperature – the attraction between cations and anions are strong and continuous, holding the ions tightly in a crystal lattice.  Crystalline structure – retain their crystal shape, even when ground into a fine powder. Solubility in water – cations are strongly attracted to the negative end of water molecules and anions are strongly attracted to the positive end of water molecules. Form electrolytic solutions – dissociate into ions; the greater the concentration of ions, the greater the conductivity. Molecular Elements and Compounds Covalent bond: forms when atoms share a pair of valence electrons so that each atom has a stable octet, resulting in the simultaneous attraction of nuclei for a shared pair of electrons. Molecule: an independent unit made up of fixed numbers of non-metallic atoms held together by covalent bonds – e.g. H2, CO2, CH4, PO43-. Molecular elements: do not exist naturally as single atoms Diatomic elements are composed of two identical atoms – H2, N2, O2, F2, Cl2, Br2, I2. Polyatomic elements are composed of many identical atoms – P4, S8, O3.  Prefix Number Rules for naming binary molecular mono- 1 compounds that do not contain di- 2 hydrogen: tri- 3 1. Name the first element. tetra- 4 penta- 5 2. Name the second element and change the ending to “-ide”. hexa- 6 hepta- 7 3. Add prefixes to indicate the octa- 8 number of atoms. nona- 9 deca- 10 Rules for writing formulas for molecular compounds: 1. Write the formula for the first element. 2. Write the formula for the second element. 3. Add subscripts to indicate the number of atoms. Example Problem 19: phosphorous tetrachloride Example Problem 20: BF3 Example Problem 21: CS2 Example Problem 22: tetraphosphorous decoxide Most molecular compounds that contain hydrogen have common names that must be memorized: 1. ammonia NH3(g) 2. glucose C6H12O6(s) 3. hydrogen peroxide H2O2(l) 4. sucrose C12H22O11(s) 5. methane CH4(g) 6. ethane C2H6(g) 7. propane C3H8(g) 8. methanol CH3OH(l) 9. ethanol C2H5OH(l) 10. water H2O(l) Properties of molecular compounds include: 1. Form non-electrolytic solutions – molecular compounds do not dissociate into ions. 2. Low solubility – most are not soluble in water. 3. Relatively low melting and boiling points –tend to have weak intermolecular forces (between molecules). 4. Require large amounts of energy to decompose – indicates that covalent bonds (intramolecular forces) within the molecule are strong. 5. Molecular substances can form crystals like ionic compounds, but they crumble easily. (2R,3R,4S,5S,6R)-2- {[(2S,3S,4S,5R)-3,4-Dihydroxy- 2,5-bis(hydroxymethyl)oxolan-2- yl]oxy}-6-(hydroxymethyl)oxane- 3,4,5-triol Water is a polar molecule – electrons are shared unequally, resulting in a slightly positive end and a slightly negative end. The positive and negative ends of neighboring water molecules are attracted to one another, forming hydrogen bonds. The polarity of water and presence of hydrogen bonding results in water having many unique properties: High melting point and boiling point – it takes a large amount of energy to separate individual water molecules from each other. Ability to absorb and release large amounts of thermal energy with only slight temperature changes.  Ice floats on liquid water – forms a six-sided crystal that is less dense than liquid water.  High surface tension – attractive force between surface molecules.  Concave meniscus and capillary action – water has cohesive and adhesive properties.  Universal solvent – positive and negative ends of water molecules are attracted to other substances. States of Matter and Solubility States of matter are communicated in brackets after the element or compound. Solid (s) Liquid (l) Gas (g) Aqueous (aq) – means dissolved in water Determining states of matter: Elements – communicated in the element box on the periodic table. Molecular compounds – generally smaller molecules tend to be gases and larger molecules tend to be liquids or solids. Ionic compounds – solid at room temperature and pressure. Ionic compounds in an aqueous environment – refer to the solubility table oIf the compound is very soluble then the state is aqueous. oIf the compound is only slightly soluble then the state is solid. Acids and Bases pH: a measure of the concentration of hydrogen ions in a solution. Every decrease of 1 on the pH scale indicates a ten- fold increase in hydrogen ion concentration (acidity).  Acid-base indicators: chemicals that change colour depending on the pH of the solution – e.g. litmus paper, phenolphthalein. The following chart summarizes the properties of acids and bases: Acids Bases Taste sour Taste bitter React with metals to Feel slippery produces hydrogen gas pH less than 7 pH greater than 7 Conduct electricity Conduct electricity Turn blue litmus red Turn red litmus blue Neutralized by bases Neutralized by acids Acids are molecular compounds that ionize in water to release hydrogen ions. To recognize an acid, look for the hydrogen symbol on the left or –COOH on the right. Compounds only become acidic when they are in an aqueous environment. When the state of matter is solid, liquid, or gas, follow ionic compound naming rules The following table summarizes the naming rules for acids once they are in an aqueous environment. The following table summarizes the naming rules for acids. Substance Name Acid Name hydrogen ________ide hydro_________ic acid eg. hydrogen chloride hydrochloric acid HCl(aq) HCl(g) hydrogen ________ate _____________ic acid eg. hydrogen chlorate chloric acid HClO3(aq) HClO3(s) hydrogen ________ite _____________ous acid eg. hydrogen chlorite chlorous acid HClO2(aq) HClO2(s) Bases are usually ionic hydroxides that dissociate in water to release hydroxide ions. The presence of a hydroxide ion (OH-) with a metal ion or an ammonium ion usually indicates that a substance is basic. Compounds only become basic in an aqueous environment. Example Problem 23: H2S (g) is dissolved in water. Give the name of the acid. Example Problem 24: What is the chemical formula for sulfurous acid? Example Problem 25: H2SO4 (s) is dissolved in water. Give the name of the acid. Example Problem 26: provide the name for Mg(OH)2(s). Is this an acid or a base? Chapter 3: Chemical Change Characteristics of Chemical Change Chemical change: occurs when a substance or substances – reactants – react to create a new substance or substances – products – with new properties. Physical change: occurs when a substance undergoes a change in state. Common characteristics of chemical reactions include: 1. The production of new substances with their own characteristic properties (e.g. state, melting point, colour, and density). 2. Energy changes such as temperature change, light, or sound. In exothermic reactions, less energy is required to break bonds than is released when new bonds form – e.g. cellular respiration and combustion. In endothermic reactions, more energy is required to break bonds than is released when new bonds form – e.g. photosynthesis and instant cold packs. Note that energy is conserved, it just changes form. 3. Changes in state – e.g. formation of a gas or a precipitate (solid). 4. Consistent with the law of conservation of mass – i.e. the mass of the reactants is equal to the mass of the products. The following evidence can help you determine if a chemical change has taken place: 1. Change in odour 2. Change in colour 3. Change in energy (temperature, light) 4. Formation of a precipitate (solid) 5. Formation of a gas (bubbles) Writing Chemical Equations Word equation: reactants and products are represented with words – e.g. hydrogen + oxygen → water. Skeleton equation: shows the formulas and states of matter for all reactants and products – e.g. H2(g) + O2(g) → H2O(l). Balanced formula equation: coefficients show that there is the same number of atoms of each type of element on the reactants and products side – consistent with Lavoisier’s law of conservation of mass – e.g. 2H2(g) + O2(g) → 2H2O(l). TED Ed – Law of Conservation of Mass Example Problem 27: Balance the following chemical equation. Li(s) + O2(g) → Li2O(s) Example Problem 28: Balance the following chemical equation. Ca(NO3)2(aq) + NaOH(aq) → Ca(OH)2(s) + NaNO3(aq) Steps for writing formula equations: 1. Identify the type of reaction – formation, decomposition, hydrocarbon combustion, single replacement, or double replacement. 2. Write the correct formulas for the reactants, including states of matter. 3. Write the correct formulas for the products, including states of mater. 4. Use coefficients to balance the equation – DO NOT CHANGE SUBSCRIPTS!! 5. Ensure coefficients represent the lowest whole number ratio of substances. Most chemical reactions can be classified as one of five main types: 1. Formation/Synthesis: two elements combine to form a compound (A + B → AB). Example Problem 29: sodium reacts with chlorine. Example Problem 30: magnesium reacts with sulfur. 2. Decomposition: a compound breaks into its elements (AB → A + B). Example Problem 31: water decomposes. Example Problem 32: decomposition of magnesium chloride. 3. Single replacement: an element reacts with a compound to produce a different compound and element (A + BC → B + AC). Example Problem 33: iron reacts with silver nitrate solution. Example Problem 34: silver bromide solution reacts with chlorine gas. 4. Double replacement: usually occurs between two ionic compounds in solution to produce two different ionic compounds (AB + CD → AD + CB). Example Problem 35: lead (II) nitrate solution reacts with sodium iodide solution. Example Problem 36: lithium hydroxide solution reacts with sulfuric acid. 5. Hydrocarbon combustion: substance containing hydrogen and carbon reacts with oxygen to produce carbon dioxide and water (CxHx + O2(g) → CO2(g) + H2O(g)). Example Problem 37: the combustion of methane. Example Problem 38: combustion of ethane. The Mole TED Ed – The Mole Mole (mol): a quantity that chemists use to measure amounts of elements and compounds. The number of entities (atoms, molecules, formula units, or ions) in 1 mole is called Avogadro’s number (NA). Avogadro’s number = 6.02 x 1023 or 602 000 000 000 000 000 000 000.  The coefficients of a balanced chemical equation represent the number of moles of each substance required for the reaction. Atomic molar mass: mass one mole of a particular element – found on the periodic table. Molar mass: the mass in grams of 1 mol of a substance – measured in grams/mol. Determined using the chemical formula of the substance and the atomic molar mass values from the periodic table. Example Problem 39: find the molar mass of glucose. Example Problem 40: calculate the molar mass of solid sulfur. Example Problem 41: how many atoms are in a 4.5 mol sample of potassium metal? The number of moles of a substance is related to its molar mass by the following equation: m=nxM m = mass (g) n = quantity of matter (mol) M = molar mass (g/mol) Example Problem 42: calculate the mass of 3.00 mol of carbon. Example Problem 43: how many moles of barium nitrate are in a 56.18 g sample?