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Chapter 2 – Atoms and Elements What is chemistry? It is often defined as “the study of matter”. It answers the questions: • “What is a substance made of?” • “How was it made?” • “How will it interact with other substances?” e.g. The chemistry of beer Beer is a homogeneous mixture consisting of water (_____), ethanol (___________), carbon dioxide (_____) and a variety of other substances responsible for its flavour. Beer is made in a multi-step process:1 1. Barley mash is heated with water. This activates enzymes in the barley that break the starch down to glucose. 2. The barley husks are filtered out of the resulting sugary water (the “wort”) which is then boiled with hops to impart flavour (by dissolving some of the more flavourful molecules from the hops). 3. The hops are filtered out, and yeast is added for the fermentation step in which they convert glucose (______) into carbon dioxide and ethanol: C6H12O6(aq) 2 CO2(g) + 2 CH3CH2OH(aq) 4. After fermentation is complete, the yeast is filtered out. The beer is then aged in tanks and filtered again before packaging. 1 www.sleeman.ca, visited June 6, 2005 How does beer interact with other substances? If certain bacteria get into the beer, their enzymes will convert the ethanol in beer to acetic acid (___________): CH3CH2OH(aq) + O2(g) CH3CO2H(aq) + H2O(l) The interactions between beer and the human body are well known (taste, inebriation, etc.) The taste is due to the shapes of the flavour molecules and how they fit into receptor molecules in our taste buds. Due to their structure, the ethanol molecules travel easily through the human body (they are soluble in both water and fat) until they reach their target, the brain. Chemistry is often termed “the central science”. It lies between biology and physics, and chemical knowledge is also necessary to truly understand many concepts in physics, biology, medicine, geology, environmental science, etc. In many ways, learning chemistry is like learning a language. The chemist’s alphabet is the collection of elements – which are combined to make compounds (“words”) – which are, in turn, combined in reaction equations (“sentences”). Just as there are rules for writing sentences that make sense, we will see that there are rules for writing reaction equations that make sense. Atomic Theory Ancient Greek philosophers proposed that all matter consisted of some combination of four elements: air, earth, fire, water. Democritus (~460-370 B.C.) disagreed, proposing that all matter could be repeatedly subdivided until an indivisible particle was reached. He called this the atom (Greek: a = not; tomos = cut). In 1785, Antoine Lavoisier (1743-1794) is credited with discovery of the law of conservation of mass: Mass is neither created nor destroyed in a chemical change. In 1794, Joseph Proust (1754-1826) demonstrated the law of definite proportions (aka law of constant composition): In a given chemical compound, the proportions by mass of the elements that compose it are fixed, independent of the origin of the compound or its mode of preparation. In 1808, John Dalton (1766-1844) based his atomic theory of matter on the ideas of Democritus, Lavoisier and Proust. 1. All matter consists of solid and indivisible atoms. 2. All of the atoms of a given chemical element are identical in mass and in all other properties. 3. Different elements have different kinds of atoms; these atoms differ in mass from element to element. 4. Atoms are indestructible and retain their identity in all chemical reactions. 5. The formation of a compound from its elements occurs through the combination of atoms of unlike elements in small whole-number ratios. While the essence of this theory has withstood the test of time, most of the postulates have since been modified: 1. Atoms can be further divided into subatomic particles. 2. Different isotopes of an element have different masses. (e.g. Carbon-12, carbon-13 and carbon-14 have masses of 12, 13.003 and 14.003 u respectively.) 3. Still true, but some have very similar masses. (e.g. Nitrogen-14 and carbon-14 both have masses of 14.003 u.) 4. In nuclear reactions, atoms do not retain their identity. 5. True; however, Dalton was unaware that not all elements are made up of single atoms. e.g. Nitrogen does not exist as free atoms; rather, two nitrogen atoms (N) bond to make one nitrogen molecule (N2). A molecule is a grouping of two or more atoms bonded together by strong attractive forces. A molecule is a discrete entity. Cl Cl Na Na O H H H Na H Cl Cl Na Cl Na Na Cl molecules Cl Na Na Cl Cl Na NaCl does not exist as discrete molecules Elements can exist as: • atoms (e.g. He, Ne) • molecules (e.g. H2, S8) • infinite materials (e.g. most metals and metalloids, carbon as diamond or graphite). Compounds can exist as: • molecules (e.g. H2O, CO, CO2) • infinite materials (e.g. quartz consists of silicon and oxygen in a 1 : 2 ratio) • simple ions (e.g. NaCl consists of Na+ and Cl-) • complex ions (e.g. CuSO4 consists of Cu2+ and SO42-). Compounds and elements can be described by chemical and physical properties: A _____________________ is one that can be observed without changing a compound/element into another compound/element. A _____________________ is one that can only be observed by changing a compound/element into another compound/element. Define the following properties as chemical or physical: 1. Water has a density of 1.000 g/mL at 4˚C. 2. Sodium reacts with water to produce sodium hydroxide. 3. Gasoline burns in the presence of oxygen to give carbon dioxide and water vapour. 4. Solid iodine sublimes to give dark purple iodine vapour. …back to Atomic Theory In 1897, Sir John Joseph Thomson (1856-1940) discovered the electron.2 He also proposed that atoms contained electrons. He reasoned that, since electrons have negative charge and atoms are neutral, atoms must also contain some constituent with positive charge. His proposal was the ‘plum pudding’ model of the atom in which electrons were embedded in a positively charged sphere. 2 previously proposed and named by physicist G. B. Stoney in 1874 + + - - + - + - + + + + - 'plum pudding' model of the atom While Thomson was able to measure the charge-to-mass ratio for the electron, Robert Andrews Millikan (1868-1953) was the first to measure the charge independently (-1.602176 × 10-19 C). This also allowed calculation of the mass (9.109383 × 10-28 g). As a matter of convenience, chemists refer to charge relative to the charge of an electron rather than in Coulombs (C). By this method, the electron has a charge of -1. In 1911, Ernest Rutherford (1871-1937) proposed a new model of the atom. He found that most alpha nucleus (charge of +78) particles fired at a thin gold foil passed 78 electrons through the foil as though traveling through empty space; however, a few were deflected by large angles. This was inconsistent with Thomson’s model. Rutherford’s ‘nuclear’ model consisted of a tiny positively charged nucleus containing most of the atom’s mass. nuclear model of a gold atom In 1919, Rutherford reported discovery of the proton, demonstrating that the nucleus of a hydrogen atom was the fundamental unit of positive charge and proposing that the nuclei of heavier atoms consisted of protons and neutral particles of similar mass. These neutral particles, ___________, were discovered by James Chadwick (1891-1974) in 1932. Thus, atoms consist of three types of subatomic particles: mass charge location outer electron 9.109383 × 10-28 g 0.0005485799 u -1 region -24 proton 1.672622 × 10 g 1.007276 u +1 nucleus -24 neutron 1.674927 × 10 g 1.008665 u 0 nucleus Defining an Element Given the extremely small masses of atoms and subatomic particles, chemists invented a new unit of mass. The atomic mass unit (u) is defined as one twelfth of the mass of a carbon atom containing six protons, six neutrons and six electrons: 1 u = 1.661 × 10-24 g As such, the mass of an atom in u will be approximately equal to the combined number of protons and neutrons it contains. Every atom has an atomic number and a mass number: mass number symbol atomic number 12 6 C Atomic number (Z) = # protons Mass number (A) = # protons + # neutrons The charge of an atom can be found by comparing the number of protons and electrons: charge = # protons - # electrons A neutral atom has an equal number of protons and electrons. The number of protons defines what element an atom belongs to. If the number of protons changes, it is not the same element. As such, writing atomic number is optional because the element symbol tells us what it must be. Mass numbers are not optional. How many protons, neutrons and electrons are in the neutral element with Z = 26 and A = 56? What element is this? Isotopes Different atoms of the same element can have different numbers of neutrons. They will have the same ____________ number but different _________ numbers. e.g. 1 2 1 1 12 13 H 6 C 6 H 3 C 14 1 6 H C These are called isotopes of the element. Only a few elements have just one naturally occurring isotope (e.g. 19F, 31P). Most elements occur as mixtures of several isotopes. Chemists normally treat these elements as consisting of “averaged” atoms with a “averaged” masses. Atomic mass (as shown on the periodic table) is the weighted average of all naturally occurring isotopes of an element. It factors in the mass and percent abundance of each isotope where % abundance = # atoms of isotope × 100% total # atoms of element atomic mass = %abund. iso.#1 %abund. iso.#2 (mass iso.#2) (mass iso.#1) + 100% 100% e.g. Chlorine has two naturally occurring isotopes: 35 37 17 17 Cl 75.8% 34.97 u Calculate the atomic mass of chlorine. Cl 24.2% 36.97 u e.g. Gallium has two naturally occurring isotopes and an atomic mass of 69.723 u: 69 71 31 31 Ga Ga 68.926 u 70.925 u Calculate the percent abundance of each isotope of gallium. Thus, we can calculate the mass of an atom. By adding atomic masses, we can also calculate the mass of a molecule (or formula unit of an infinite material). e.g. Sulphur has an atomic mass of 32.066 u. Oxygen has an atomic mass of 15.999 u. Calculate the molecular mass of SO3. In the lab, however, chemists don’t often work with single atoms or molecules. It is far more common to work with quantities in the 1 mg to 1 kg range. Recall that 1 u = 1.661 × 10-24 g. Thus, 1 g = 6.022 × 1023 u. This means that 1 gram contains a *lot* of atoms or molecules. Since the human brain has trouble comprehending numbers this large, another unit was created to make such quantities easier to discuss: 1 mole = 6.022 × 1023 Thus, 1 g = 1 mole µ Or, 1 µ = 1 g/mole This ‘magic’ number (6.022 × 1023) is called _______________ _______________ in honour of Amedeo Avogadro (1776-1856) who was the first to propose that such a number could exist. For convenience, the word ‘mole’ is often abbreviated to ‘mol’ when used as a unit. e.g. The mass of carbon-12 is 12 g/mol. The key to working with moles is to remember that 1 mole is always equal to 6.022 × 1023 of whatever type of particle you are discussing. e.g. There are __________ molecules of CO2 in 1 mole of CO2. There are __________ atoms of C in 1 mole of CO2. There are __________ atoms of O in 1 mole of CO2. How many moles of C2H6O contain 5.0 x 1024 atoms of H? How many molecules are there in 392.3 g sulfuric acid (H2SO4)? What is the mass of a sample of hydrochloric acid (HCl) containing 2.01 × 1024 atoms of hydrogen? What is the mass of a sample of nitric acid (HNO3) containing 3.011 × 1022 atoms of oxygen? The Elements Currently, ____ elements are known. Of these, ____ occur naturally while the other ____ have been made in laboratories. The Gases ____ elements are gases at room temperature The “noble gases” or “inert gases” are: Helium, Neon, Argon, Krypton, Xenon, Radon Air is mostly Nitrogen (___) and Oxygen (___) (~80% N2 + ~20% O2) The other gases are: Fluorine (symbol ___, formula ____, highly reactive) Chlorine (symbol ___, formula ____, reactive) Hydrogen (symbol __, formula ___, flammable) The Liquids ____ elements are liquids at room temperature: Bromine (symbol ____, formula ____, corrosive) Mercury (symbol ____, formula ____, poisonous) The Solids The remaining elements are solids at room temperature. Chemists have a special tool for recognizing which elements are similar and which are different. This tool is known as the ________________________. The Periodic Table In 1869, Dmitri Mendeleev (1834-1907) noticed that certain elements exhibited similar behaviour – most notably, the ratios with which they formed molecules with hydrogen and with oxygen. By arranging the elements in order of increasing mass and such that similar elements formed columns, he developed the first periodic table: This periodic table was incomplete – all of the inert gases are missing, but it was remarkably accurate in other respects. If there appeared to be a ‘missing’ element, Mendeleev left a blank space, assuming that it would be discovered at a later date. He was proven correct with the discoveries of _________ (69.7 u) in 1875 and _______________ (72.6 u) in 1886. In 1913, H.G.J. Moseley (1887-1915) noted that the periodic table would be more descriptive if the elements were listed in order of increasing _____________ rather than increasing mass. This led to the modern periodic table and law of periodicity: “The properties of the elements are periodic functions of atomic number.” We will see why this is after looking at some of the properties of the elements. In the modern periodic table, a column is known as a ________________ and a row is known as a _______________. Elements in the same period have similar ___________ but very different ________________. Elements in the same group have similar _________________ but very different ____________. Despite having similar chemical properties, elements in the same group sometimes have quite different physical properties. Why might this be? main groups The periodic table can also be divided to give different classes of elements. The upper right portion of the periodic table consists of nonmetals while the lower left portion consists of metals and a few elements at the border are metalloids. main groups transition groups A Quick Overview of the Elements by Group Hydrogen (H) • shares similarities with elements in groups 1 and 17 but doesn’t fully belong to either group • diatomic gas (H2) that is fairly unreactive • reacts explosively with oxygen to make water (H2O) if enough energy is supplied • three isotopes: 1H is the most common (and the only element with no neutrons), 2H is in heavy water used to cool nuclear reactors; 3H is radioactive Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr) • aka Group 1A • soft metals that react strongly with water to give XOH and with oxygen to give X2O (reactivity increases with atomic mass) • do not exist in pure form in nature due to high reactivity • readily lose 1 electron to make cations with +1 charge Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra) • aka Group 2A • most are metals that react with water to give X(OH)2 and with oxygen to give XO (reactivity increases with atomic mass; beryllium does not react with water) • do not exist in pure form in nature due to high reactivity (though react less strongly with water than the Group 1 element in the same period) • beryllium is highly toxic • readily lose 2 electrons to make cations with +2 charge Groups 3-12: Transition Metals • aka Groups 1B-8B (but not in order!) • Group 3 metals readily lose 3 electrons to make cations with +3 charge; otherwise, they act like Group 2 metals • Groups 4-11 are the ‘true’ transition metals in that they lose electrons to form coloured compounds in which the metal atom has a positive charge • Groups 11 and 12 mimic the behaviour of Groups 1 and 2 respectively but are less reactive • the metals in the middle of the transition groups are hardest (especially tungsten (W) in Group 6) • the metals at the edges of the transition groups (i.e. Groups 3 and 12) are the softest • mercury is the only liquid metal; the rest are solids • copper and gold are the only coloured metals • silver is the best conductor of electricity • gold is the most malleable metal • readily lose electrons to make cations Group 13 (B, Al, Ga, In, Tl) • aka Group 3A • most are metals; boron is a metalloid • all are solids; however, gallium has a very low melting point (29.76 ˚C) • aluminum is an industrially important lightweight metal and the third most abundant element in the earth’s crust • compounds containing indium and thallium are toxic • form compounds in a 1:3 ratio with halogens (e.g. BCl3) • lose 3 electrons to make cations with +3 charge Group 14 (C, Si, Ge, Sn, Pb) • aka Group 4A • carbon is a nonmetal; silicon and germanium are metalloids; tin and lead are metals • carbon exists in several different allotropes (graphite, diamonds, fullerenes), and compounds containing carbon form the basis for life • silicon is the second most abundant element in the earth’s crust; it does not naturally occur in pure form but as silicates (compounds made of silicon and oxygen) which form rocks, sand, glass, etc. • form compounds in a 1:4 ratio with halogens (e.g. CCl4) Group 15: Pnictogens (N, P, As, Sb, Bi) • aka Group 5A • nitrogen and phosphorus are nonmetals; arsenic and antimony are metalloids; bismuth is somewhat metallic • nitrogen is a highly stable diatomic gas (N2) and the most abundant element in the atmosphere • phosphorus exists in three allotropes (white phosphorus is P4; red and black phosphorus are polymer used in match heads) • form compounds in a 1:3 ratio with hydrogen (e.g. NH3) Group 16: Chalcogens (O, S, Se, Te, Po) • aka Group 6A • oxygen, sulphur and selenium are nonmetals; tellurium is a metalloid; polonium is a metal • oxygen is the most abundant element in the earth’s crust & the second most abundant element in the atmosphere • oxygen exists in two allotropes, O2 (oxygen) and O3 (ozone); both are reactive gases • sulphur exists in many allotropes (S2, S6, S8, etc.) • except for oxygen, chalcogens tend to make compounds with unpleasant odours (which get worse as atomic mass increases) • form compounds in a 1:2 ratio with hydrogen (e.g. H2O) • gain 2 electrons to make anions with -2 charge Group 17: Halogens (F, Cl, Br, I, At) • aka Group 7A • nonmetals that exist as diatomic molecules (except for astatine which is too unstable to study) • fluorine and chlorine are gases; bromine is a liquid; iodine is a solid • colourful (F2 is pale yellow; Cl2 is yellow-green; Br2 is red-brown; I2 is dark purple) • form compounds in a 1:1 ratio with hydrogen (e.g. HF) • gain 1 electron to make anions with -1 charge Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn) • aka Group 8A, inert gases or rare gases • unreactive gaseous nonmetals (a few compounds have been made containing Xe and Kr but they are rare) • helium is used to fill blimps, etc. due to low density • 1% of the atmosphere is argon, and argon is often used as an inert atmosphere in labs because it is denser than air and unreactive Important Concepts from Chapter 2 • chemical vs. physical properties • law of conservation of mass • Dalton’s atomic theory of matter • Thomson and Rutherford models of the atom • subatomic particles (protons, neutrons, electrons) • atoms, molecules, infinite materials, simple and complex ions • periodic table (groups and periods) • elements (names and symbols) • atomic number and mass number • isotopes • calculating average atomic mass and percent abundance • Avogadro’s number and the mole