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AP Chemistry Chapter 2 Atoms, Ions, and Nomenclature
History of Atomic Theory Democritus (p. 40) - supported the idea that matter is continuous, infinitely divisible into smaller pieces, or composed of small, indivisible particles. Lavoisier- (p. 41) father of modern chemistry; Law of Conservation of Mass (mass is neither created nor destroyed in ordinary chemical reactions) Joseph Proust (p. 42) – Law of Definite Proportions (a given compound always contains exactly the same proportion of elements by mass) (For example: water is always made up of two hydrogen atoms and one oxygen atom. Table salt is always made up of one sodium atom and one chlorine atom.) John Dalton-modern atomic theory (p ); Law of Multiple Proportions (when new elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to whole numbers)(For example: in H 2 O, 2 g H will combine with 16 g O (1:8 ratio), but in H 2 O 2, 2 g H will combine with 32 g O (1:16 ratio))(see sample ex. 2.1, p. 43)
2.3 Dalton’s Atomic Theory Elements are made from tiny particles called atoms. All atoms of a given element are identical. (see isotopes below) The atoms of a given element are different from those of any other element. Atoms of different elements combine to form compounds. A given compound always has the same relative numbers and types of atoms. Atoms cannot be created or destroyed in an ordinary chemical reaction; they are simply rearranged to form new compounds. (Law of Conservation of Mass)
2.4 Early Experiments to Characterize the Atom ScientistExperimentKnowledge GainedRelating to CrookesCathode Ray Tube Negative particles of some kind exist electron J.J. Thomson Late 1890’s Fig. 2.8, p. 49 Cathode Ray deflection Mass/charge ratio of the electron determined e/m = x 10 8 C/g Electron, fig. 2.9 Robert Millikan 1909, fig. 2.10, p. 50 Oil drop experiment Mass of the electronElectron Ernst Rutherford, Marsden, Geiger Fig. 2.11, p. 51 Gold foil experiment, 2.12 Nucleus present in atom Fig Nucleus of an atom proton
James Chadwick- discovered neutron (not in textbook) Niels Bohr- proposed idea that the atom was made up of the nucleus containing protons & neutrons that was being orbited by electrons in specific allowed orbits. This particle model of the electron and atom was expanded a few years later after Bohr’s original ideas to incorporate the wave nature of the electrons. (textbook, chap. 11)
2.5 Modern View of Atomic Structure ParticleChargeMass, in kilogramsMassLocation Proton x amuNucleus Neutron01.67 x amuNucleus electron9.11 x /1836 amu = 0 amu (approximately) Electron cloud outside nucleus
Z = atomic number = number of protons in the nucleus of one atom of an element (p. 54) (Because all atoms are neutral, it also tells us how many electrons are surrounding the nucleus. When atoms gain or lose electrons, they become charged and form ions.) A = mass number = the number of protons + the number of neutrons in one atom of the element (p. 54) Example: Sr Hg Problem: Magnesium has an atomic number of 12. Write the symbol identity for an atom of magnesium with a mass number of 24. ( Mg) Write the symbol for silver (Z = 47) that has 61 neutrons. ( Ag) Give the symbol for a phosphorus atom (Z = 15) that contains 17 neutrons. ( P)
Isotopes Atoms with the same number of protons & electrons, but different numbers of neutrons are called isotopes. For example, hydrogen has three isotopes Protium 1 1 HDeuterium 2 1 H Tritium 3 1 H There are two isotopes of carbon, carbon-12 ( 12 6 C) and carbon-14 ( 14 6 C).Since it is the electrons in atoms that affect the chemical properties of a substance, isotopes of the same element have the same chemical properties.
2.6 Molecules and Ions Molecules are formed when a definite arrangement of atoms are joined together by chemical bonds. A molecule can consist of the atoms of only one element or the atoms of many different elements, but always in a fixed proportion. This means that molecules can be elements or compounds. Molecules are usually formed between non-metal elements. See table in notes
Ions Atoms have equal numbers of protons and electrons and consequently have no overall charge. When atoms gain or lose electrons, the proton-electron numbers are unbalanced, causing the particles to become charged. These charged particles are called ions. Positive ions (where # protons > # of electrons) are called cations; negative ions (where # protons < # electrons) are called anions. Metals tend to form cations and non-metals tend to form anions. Oppositely charged ions will join together to make ionic compounds. An ion made up of only one type of atom is called a monatomic ion; one made up from more than one type of atom is called a polyatomic ion. See table in notes.
2.7 The Periodic Table Vertical columns are called groups. Elements in the same group have similar properties. Some groups have special names: 1- alkali metals17- halogens 2- alkaline earth metals18- noble gases Horizontal rows are called periods. The lanthanide series and the actinide series are technically part of periods 6 & 7, respectively. Depending on the author of the particular periodic table, these two series may differ in their elements. Metals are good conductors of heat and electricity, are malleable, ductile (able to be drawn into a fine wire) and are usually shiny (have luster). (All metals, except mercury, are solids at room temperature and pressure.) Non-metals generally lack the properties of the metals and show much more variation in their properties; some are solids, bromine is a liquid, and many are gases at room temperature and pressure. Metalloids (semimetals) exhibit properties common to both metals and non-metals. For example, silicon conducts electricity, but has a higher melting point than most metals.
2.8 Naming Simple Compounds (Organic compounds are carbon-containing compounds and have their own nomenclature which we will study in Unit 17.) Binary Ionic Compounds (Type I) Binary compounds are formed between two elements. In compounds where one is a metal and one is a non-metal, an ionic compound is formed. The formulas and names can be determined by considering the charge on the ions. To find the formula of an ionic compound, the positive and negative charges must be balanced, i.e., there must be no net charge. The unmodified name of the positive ion is written first, followed by the root of the negative ion with an –ide suffix.
Binary Ionic Compounds (Type II) Most transition metal ions include a number written as a Roman numeral after the name. Most transition metals have varying charges and the number identifies the charge on the ion. Older nomenclature gave special suffixes to the different numbers. The higher number ends in –ic; the lower one, in – ous. CuClcopper (I) chloride CuCl 2 copper (II) chloride SnOtin (II) oxide SnO 2 tin (IV) oxide
Binary Covalent Compounds (Type III) A compound made from 2 elements that are both non-metals, then the compound is molecular. The name of the first element is written first and the name of the second element has the suffix -ide added to the root of the element name. Sometimes two elements can combine in different ratios. Prefixes are then added to the element names to distinguish these possibilities. The prefix mon- is never to the first element if only one atom is present. If the prefix ends in a or o and the name of the element begins with an a or an o, then the final vowel of the prefix is omitted. Some compounds, such as water, are not named using this system. See table in notes for prefixes.
NOnitrogen monoxide NO 2 nitrogen dioxide N 2 O 5 dinitrogen pentoxide P 2 O 5 diphosphorus pentoxide P 4 O 10 tetraphosphorus decoxide
Compounds containing Polyatomic Ions (p. 66) Polyatomic ions are those where more than one element makes up a unit with an overall ionic charge. Polyatomic ions that contain oxygen and another non-metal are called oxyanions and can be named systematically. In these oxyanions certain non-metals (Cl, N, P, and S) form a series of oxyanions containing different numbers of oxygen atoms. See table in notes for examples.
Criss-cross method of writing formulas Write the positive charged ion first; then write the negatively charged ion Take the number of the charge (but not the sign) on the anion and write it as the subscript on the cation. Take the number of the charge (but not the sign) on the cation and write it as the subscript on the anion. Make sure the numbers are reduced to the lowest ratio (for example, Mg 2 O 2 reduces to MgO) You cannot reduce a subscript which already exist in a polyatomic formula, for example, peroxides: Na 2 O 2 does not reduce to NaO)
When using the criss-cross method to write the formula of a compound containing one or two polyatomic ions, a set of parentheses will surround any polyatomic ion for which a number other than one (1) is needed as a subscript. This does not include the original subscript which may exist in the polyatomic ion’s formula. If only two forms of a polyatomic ion exist, the suffix endings are –ite (for the fewer number of oxygen atoms) and –ate (for the larger number of oxygen atoms) Some polyatomic ions contain hydrogen and are named accordingly. The prefix thio- means that a sulfur atom has replaced an atom of oxygen in an anion.
Na 2 CO 3 sodium carbonate NH 4 NO 3 ammonium nitrate Mg 3 (PO 4 ) 2 magnesium phosphate CaSO 4 calcium sulfate
Acids (p. 71) In this consideration of nomenclature, an acid is a compound that produces hydrogen ions when it is dissolved in water. A binary acid is a compound that contains an anion with no oxygen present. Such compounds are named by placing the prefix hydro- followed by the non-metal name modified by an –ide suffix.
Oxyacids (p. 71) Oxyacids are formed when hydrogen ions combine with oxyanions. This makes a ternary compounds (one containing three elements: hydrogen, oxygen, and another non-metal). If the anion ends in –ite, then the oxyacid ends in –ous; if the oxyanion ends in –ate, then the acid will end in –ic. H 2 SO 3 sulfurous acid H 2 SO 4 sulfuric acid
Hydrates (not in textbook) Hydrated compounds contain water molecules incorporated into the molecular structure. Strong heating can generally drive off the water molecules from these ionic compounds (salts). (Occasionally, heating will cause the compound to decompose, not just simply dry out.) Once the water has been removed, the salt is said to be anhydrous (without water). CuSO 4. 5H 2 Ocopper (II) sulfate pentahydrate