Chem Lec Reviewer (Atoms, Molecules, Ions, Nomenclature) PDF

Dalton’s Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century (1808), championed by John Dalton. Democritus (400 BC) – ancient Greek philosopher, credited with being one of the earliest fig...

Dalton’s Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century (1808), championed by John Dalton. Democritus (400 BC) – ancient Greek philosopher, credited with being one of the earliest figures to propose the idea of the atom. DALTON’S POSTULATES 1. Each element is composed of extremely small particles called atoms. (“atomos”) 2. All atoms of a given element are identical (the same size, weight, and form), but the atoms of one element are different from the atoms of all other elements. 3. Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions 4. The compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atom. Dalton developed the atomic theory to explain: 1. The Law of Conservation of Mass 2. The Law of Definite Proportions 3. The Law of Multiple Proportions LAW OF MASS CONSERVATION - The total mass of substances does not change during a chemical reaction LAWS OF DEFINITE ( OR CONSTANT ) COMPOSITION no matter what its source, a particular chemical compound is composed of the same elements in the same proportions by mass. LAW OF MULTIPLE PROPORTIONS If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers Revisions of the atomic structure: NOT all atoms of any given element are identical (came about by the discovery of isotopes) Isotopes are atoms of the same kind of element but have different mass number  Some atoms may be disintegrated. - ex. Radioactive elements - Radioactive elements undergo constant, spontaneous disintegration termed radioactive decay to form atoms of simpler elements ATOMS - are not as indivisible as Dalton imagined - are not the smallest particles - are composed of subatomic particles - The three most important are: Electron Proton Neutron ELECTRON - knowledge of electron began with a study of the discharge of electricity through gases at reduced pressure. - J.J Thomson was credited for the discovery of electrons - Cathode Ray Tube Experiment - Millikan Oil Drop Experiment - determine the charge on the electron in 1909 PROTON - with the discovery of electrons, the concept aroused that if there are negatively charged particles, there should also be positively charged particles - 1886, E. Goldstein discovered that there were rays travelling in the opposite direction to that travelled by the cathode rays NEUTRON - Identified by J. Chadwick in 1932 - Neutral units of matter SUBATOMIC PARTICLES Protons (+1) and electrons (-1) have a charge; neutrons are neutral Protons and neutrons have essentially the same mass (relative mass = 1). The mass of an electron is so small we ignore it (relative mass = 0) Protons and neutrons are found in the nucleus; electrons travel around the nucleus HISTORICAL DEVELOPMENT OF AN ATOMIC MODEL J.J Thomson in 1898, suggested an atom model: PLUM PUDDING MODEL An atom is composed of a sphere of positive electricity with embedded electrons. This concept was abandoned in 1911 as a result of Rutherford’s experiment. The NUCLEAR ATOM  THE NUCLEAR ATOM (Rutherford’s Experiment in 1910) use α particle to study the inner structure of atoms a beam of α-particles was directed at a thin gold foil the majority of α-particles penetrated the foil undeflected some α- particles experiences slightly deflections A few (about one in every 20,000) suffered rather serious deflections as they penetrated the foil. Interpreting Rutherford’s gold foil experiment: → An atom is spherical and contains a very dense center (the nucleus) → The heavy subatomic particles (proton) are found in the nucleus → The very light (electron) are found outside the nucleus An atom is almost an empty space. It has a massive “core” that contains the mass of the atom. (1911) Modern View of an Atomic Structure Size of an atom Given: diameter of uranium atom = 2.5 x 10-8 cm diameter of a coin = 1.9 cm Required: How many uranium atoms placed side by side would it take to span the diameter of the coin? Answer: 76 million uranium atoms Atomic Mass Unit (amu) A unit of mass, use for atoms 1 amu = 1.66054 x 10-24 g Angstrom (A or Å) A unit of length 1 A = 1 x 10-8 cm = 1 x 10-10 m What is the diameter of uranium in angstrom 2.5 A ATOMIC NUMBER, MASS NUMBER Atomic Number (Z) - number of PROTONS in the nucleus of an atom - This is the defining trait of an element. Its value determines the identity of the atom - number of PROTONS  Mass Number (A) - a count of the total number of protons and neutrons in an atom’s nucleus - It is a whole number - number of PROTONS + number of NEUTRONS A - Z = number of NEUTRONS Atoms are ELECTRICALLY NEUTRAL, which means that they contain the same number of positive and negative charges # of protons = # of electrons NUCLIDE SYMBOL / NUCLEAR SYMBOL → Elements are presented by a one or two letter symbol → Atomic number, Z, is written as a subscript BEFORE the symbol → The mass number is written as a superscript BEFORE the symbol IONS: # of protons ≠ # of electrons When the numbers of these subatomic particles are not equal, the atom is electrically charged can either be positive charged or negatively charged The charge of an atom is defined as follows: Atomic charge = number of protons - number of electrons ISOTOPES - atoms of the same element with different masses - atoms of the same element that have different number of neutrons, but the same number of protons Atomic Mass  - atoms have extremely small masses - the heaviest known atoms have a mass of approximately 4 x 10-22 g. - A mass scale on the atomic level is used, where an atomic mass unit (amu) is the base unit - 1 amu = 1.66054 x 10-24 g - the mass of an atoms - equivalent to the sum of the number of protons and neutrons in an atom’s nucleus, since electrons have negligible mass in comparison - a weighted average of all the isotopes of an element - the sum of the atomic mass of an isotope each multiplied by the isotopes % abundance FORMULA: Atomic Mass/Weight = ∑(isotope mass) x (fractional natural abundance) ATOMIC WEIGHT MEASUREMENT atomic and molecular weight can be measured with great accuracy using a mass spectrometer ISOTOPIC ABUNDANCE - “relative abundance”, “natural abundance” - the abundance of isotopes of a chemical element as naturally found on a planet - % abundance of isotope in a sample of element was determined by mass spectrometry - the relative proportions of the stable isotopes of each element - they are most often quoted as atom percentages - the sum of the percent natural abundances of all the isotopes of any given element must total 100% FORMULA: AVOGADRO’S NUMBER AND THE MOLE to describe very large number of particles (atoms, ions, or molecules) Mole, n abbreviated as “mol”; symbol for mole, “n” an SI unit for the amount of substance 1 mole = 6.022045 x 1023 particles Note: we will round this number to 6.022 x 1023 1 mole of H atoms contains 6.022 x 1023 atoms of H MOLAR MASS, MM relating to the mass of an atom to the mass of Avogadro’s number (1 mol) of these atoms: → The mass of a single atom of an element in amu is numerically equal to its mass in grams of 1 mole of that element mass of an element in amu = mass of the element in grams mole of that element molar mass = mass in grams ; mole (n) = mass (g) mole molar mass TWO TYPES OF COMPOUNDS: IONIC COMPOUND MOLECULAR COMPOUND COMPOUNDS EMPIRICAL FORMULA gives the simplest whole number ratios of the atoms present in a compound gives only the relative number of atoms in each type of a molecule. The subscripts are always whole number ratios does not directly tell you the molecular formula, but with given information, the molecular formula can be known may or may not represent the true chemical formula of the compound MOLECULAR FORMULA represents the chemical formula for molecules or molecular compounds gives the actual number of atoms and types of atoms in a molecule provide greater information about the molecule A molecular substance is represented by a molecular formula A molecular formula tells only the number of atoms of each element in a molecule of the compound A condensed structural formula shows all atoms, but it omits the vertical bonds and most or all of the horizontal single bonds A condensed structure uses parentheses to show that polyatomic groups within a formula are attached to the nearest non-hydrogen atom on the left A structural formula uses lines to show the bonds between the atoms. A structural formula tells us which atoms are in a molecule. It also shows us how they connect to each other CATIONS - metals tend to lose electrons to form cations - positive ANIONS  - nonmetals tend to gain electrons to form anions - negative MONOTATOMIC IONS - ions that contain only one atom POLYATOMIC IONS - composed of more than one atom, but function as a unit - are negatively charged except for NH4+ , ammonium ion IONIC COMPOUNDS - compounds containing cations and anions - formed by the electrostatic attraction of cations and anions - occur in such a ratio such that the positive charges are equal to the negative charges therefore it is neutral - the cation is a metallic element and the anion is a nonmetallic element PERCENTAGE COMPOSITION - the percentage by mass of each type of element in a compound - The % by mass of each element in a compound is calculated using the equation The subscripts in the formula are first used to calculate the mass of each element in one mole of the compound The percent composition of a given compound is always the same as long as the compound is true MASS OF THE ELEMENT FROM PERCENTAGE COMPOSITION  CLASSIFICATION OF INORGANIC COMPOUNDS 1. Acids General Formula: H + a nonmetal - Binary Acids - Ternary Acids 2. Bases General Formula: a metal + OH- 3. Salts General Formula: a metal + a nonmetal - Binary Salts - Ternary Salts 4. Oxides General Formula: a nonmetal + oxygen - Metallic Oxides ( Basic Anhydrides) - Nonmetallic oxides (Acid Anhydrides 5. Binary Molecular Compounds NAMING OF COMPOUNDS…