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All matter is composed of atoms. Understanding the structure of atoms is critical to understanding the properties of matter HISTORY OF THE ATOM 1808 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS DALTONS ATOMIC THEORY Particle Mass (g) Charge (Coulombs) Electron (e-) 9.1 x 10-28 -1.6 x 10-19 Proton (p) -24 1.67 x 10 Neutron (n) 1.67 x 10-24 +1.6 x 10 0 -19 Charge (units) -1 +1 0 mass p = mass n = 1840 x mass e- HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON HISTORY OF THE ATOM 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. they fired Helium nuclei at a piece of gold foil which was only a few atoms thick. they found that although most of them passed through. About 1 in 10,000 hit Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m Atoms are composed of -protons – positively charged particles -neutrons – neutral particles -electrons – negatively charged particles Protons and neutrons are located in the nucleus. Electrons are found in orbitals surrounding the nucleus. Atomic Structure HELIUM ATOM Shell proton + electron N N + - neutron Atomic Structure Every different atom has a characteristic number of protons in the nucleus. atomic number = number of protons Atoms with the same atomic number have the same chemical properties and belong to the same element. Atomic Structure Each proton and neutron has a mass of approximately 1 dalton. The sum of protons and neutrons is the atom’s atomic mass. Isotopes – atoms of the same element that have different atomic mass numbers due to different numbers of neutrons. ATOMIC STRUCTURE Atomic mass the number of protons and neutrons in an atom 4 Atomic number the number of protons in an atom 2 He number of electrons = number of protons ATOMIC NUMBER (Z) = number of protons in nucleus MASS NUMBER (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons ISOTOPS are atoms of the same element (X) with different numbers of neutrons in the nucleus Mass Number A ZX Atomic Number 1 1H 235 92 2 1H U Element Symbol (D) 238 92 3 1H U (T) 1 1 H 2 1 H (D) 3 1 H (T) * An element can have several naturally occurring isotopes. * These isotopes of a element behave in the same way. * In calculating the relative atomic mass of an element with isotopes, the relative mass and proportion or percentage of each is taken into account. Isotope Relative isotopic mass Relative abundance (%) Cl 34.969 75.80 Cl 36.966 24.20 Ar = (relative isotopic mass X1 % abundance) + relative isotopic mass X2 % abundance) 100 Ar (Cl) = (34.969 X 75.8) + ( 36.966 X 24.2) 100 Ar (Cl) = 2650.65 + 894.58 100 Ar (Cl) = 35.45 amu (atomic mass unit) HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons. MULTIELECTRON ATOMS Quantum Number * The principal quantum number, n, describes the energy level on which the orbital resides. * Largest E difference is between E levels * The values of n are integers > 0. * 1, 2, 3,...n. Azimuthal Quantum Number, l * defines shape of the subshell * Allowed values of l are integers ranging from 0 to n − 1. Value of l Type of subshell 0 s 1 p 2 d 3 f So each of these letters corresponds to a shape of orbital. *Describes the three-dimensional orientation of the orbital. *Values are integers ranging from -l to l: −l ≤ ml ≤ l. *Therefore, on any given energy level, there can be up to: * 1 s (l=0) orbital (ml=0), * 3 p (l=1) orbitals, (ml=-1,0,1) * 5 d (l=2) orbitals, (ml=-2,-1,0,1,2) * 7 f (l=3) orbitals, (ml=-3,-2,-1,0,1,2,3) The energy of electron only depends on the value of n shell = all orbitals with the same value of n subshell = all orbitals with the same value of n and l an orbital is fully defined by three quantum numbers, n, l, and ml Each shell of QN = n contains n subshells n = 1, one subshell n= 2, two subshells, etc Each subshell of QN = l, contains 2l + 1 orbitals l = 0, 2(0) + 1 = 1 l = 1, 2(1) + 1 = 3 *Value of l = 0. *Spherical in shape. *Radius of sphere increases with increasing value of n. *Value of l = 1. *Have two lobes with a nodal plane between them. Note: always 3 p orbitals for a given n *Value of l is 2. *2 nodal planes *Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center. Note: always 5 d orbitals for a given n. *This leads to a fourth quantum number, the spin quantum number ms. *The spin quantum number has only 2 values +1/2 and -1/2 *Describes magnetic field vector of electron * *No two electrons in the same atom can have exactly the same energy. *For example, no two electrons in the same atom can have identical sets of quantum numbers. *Name of each electron unique *Name consists of four numbers: *n,l,ml,ms *Example: *Mr. George Herbert Walker Bush *We must learn to name our electrons *Unlike people, there is a lot in the “name” of an electron. *Distribution of all electrons in an atom. *Consist of * Number denoting the energy level. * Letter denoting the type of orbital. * Superscript denoting the number of electrons in those orbitals. *Each box represents one orbital. *Half-arrows represent the electrons. *The direction of the arrow represents the spin of the electron. 1s22s1 (of maximum multiplicity) NOT: “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.” *We fill orbitals in increasing order of energy. *Different blocks on the periodic table, then correspond to different types of orbitals. * Remember: The periodic table was arranged the way it was based on chemical properties. * Totally empirical, until now. Based only on observation. Short cut for writing electron configurations Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row. For instance, the electron configuration for Chromium, is [Ar] 4s1 3d5 rather than the expected [Ar] 4s2 3d4. Atomic Structure Neutral atoms have the same number of protons and electrons. Ions are charged atoms. -cations – have more protons than electrons and are positively charged -anions – have more electrons than protons and are negatively charged An ion is formed when an atom, or group of atoms, has a net positive or negative charge (why?). If a neutral atom looses one or more electrons it becomes a cation. Na 11 protons 11 electrons Na+ 11 protons 10 electrons If a neutral atom gains one or more electrons it becomes an anion. Cl 17 protons 17 electrons Cl- 17 protons 18 electrons