ATOMIC STRUCTURE HISTORY OF THE ATOM 460 Bc Democritus develops the idea of atoms -he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMOS (greek for indivisible) 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 1898 Joseph John Thomson -found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON 1904 -Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge like plums surrounded by pudding. PLUM PUDDING MODEL 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 -They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back. Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus. -He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction -However, this was not the end of the story. 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. SUMMARY 1. 2. 3. 4. 5. The Atomic Number of an atom = number of protons in the nucleus. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. The number of Protons = Number of Electrons. Electrons orbit the nucleus in shells. Each shell can only carry a set number of electrons. Atoms -came from the Greek word atomos, which means “indivisible” or “uncuttable”. Atomism -it refers to the philosophy that states that all matter was produced from atoms. -was never broadly accepted in the prehistoric world. Four Fundamental elements 1. Fire 2. Air 3. Earth 4. Water Four properties that combined to produce 4 elements 1. Hot 2. Cold 3. Dry 4. Wet Dalton’s Atomic Theory 1. Elements are made up of all particles, called atoms. 2. In any given pure element, the mass and other properties of all atoms are the same. Atoms of different elements differ in mass and other properties. 3. Compounds are composed of atoms of more than one element. The constituent atoms in a given compound are present in the same whole number ratio. 4. In a chemical reaction, atoms are neither created nor destroyed. They simply combine, separate, or rearrange. Three Laws Of Dalton 1. Law of Definite proportions -state that when a product is formed, elements always combine in similar proportions by mass regardless of the size of the sample. 2. Law of multiple proportions -this law illustrates that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. 3. Law of conservation of mass - it states that mass is neither created nor destroyed during a chemical reaction. How do we know the number of protons in an atom? • Atomic number (#)= # of protons in an atom – Ex: Hydrogen’s atomic # is 1 • hydrogen has 1 proton – Ex: Carbon’s atomic # is 6 • carbon has 6 protons **The number of protons identifies the atom-it’s an atom’s fingerprint. How do we know the number of neutrons in an atom? • Mass #: the # of protons plus neutrons in the nucleus • # of neutrons = mass # - atomic # Example – Li has a mass # of 7 and an atomic # of 3 • Protons = 3 (same as atomic #) • Neutrons= 7-3 = 4 (mass # - atomic #) How do we find the number of electrons in an atom? • Most atoms are neutral (have no overall charge) • Because the only charged subatomic particles are the protons and electrons… they must balance each other out in an electrically neutral atom. • Therefore.. – # Electrons = # Protons * * (in a neutral atom..) Isotopes • Dalton’s 1st postulate was wrong. Atoms of the same element can be different (they can have different # of neutrons) – Thus, different mass numbers. • These are called isotopes. Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. We name the isotope based on its mass number – carbon-12 – carbon-14 – uranium-235 Atomic Mass How heavy is an atom of oxygen? It depends.. b/c there are different oxygen isotopes. We are more concerned with the average atomic mass. This is determined based on the abundance of each isotope We don’t use grams for this mass because the numbers would be too small. Measuring Atomic Mass • Instead we use the Atomic Mass Unit (amu) – defined as one-twelfth the mass of a carbon-12 atom. • Each isotope has its own atomic mass, thus we determine the average from percent abundance. Quantum Model of the Atoms Louie de Broglie -a French scientist, in 1924 initially thought that electrons are behaving both like a wave and a particle. -however, later on he explained that electrons behave more of a wave-that the electrons in Bohr’s quantized orbit was similar to the observed behavior of a wave. -he proposed that electrons be considered as a wave confined in the space surrounding an atomic nucleus. Erwin Schrodinger -an Austrian physicist, in 1926, he developed an equation that treated electrons as behaving like a wave. Werner Heisenberg -a German physicist, in 1927, he tried to address the possibility that electrons are behaving both as a wave and a particle. Heisenberg’s uncertainty principle -state that it is impossible to determine simultaneously both position and momentum ( product of mass and velocity) of an electron. Quantum theory -explains the mathematical description of the wave properties of electrons and other small particles. -it suggest the probability of finding an electron at a given place around the nucleus. They exist in a certain regions called orbitals. Orbitals -is a three-dimensional region surrounding the nucleus, which indicates probable location of the electron. Quantum numbers -state that the properties of the atomic orbitals and the properties f the electrons in orbitals. Principal Quantum number (n) -indicates the size of the orbital. The bigger the n is, the greater the average distance of an electron in the orbital from the nucleus and thus, the larger the orbital and the atom. -shows the main energy level occupied by an electron, and takes positive integers (1,2,3, and so on) as values. Shell - is composed of set of orbitals that have the same principal quantum number. Angular Momentum Quantum Number (l) -also known as azimuthal quantum number. -it determines or corresponds to the shape of the orbital. Allowed values for l include 0 and all positive integers less than or equal to n-1. Subshells - are one or more orbitals with the same values of n and l. Magnetic quantum number (ml) -indicates the orientation of an orbital around the nucleus. Spin quantum number (ms) - may have only two possible values, + ½ and – ½, which indicates the spin of the electrons The Shapes of Atomic Orbitals • the l quantum number primarily determines the shape of the orbital • l can have integer values from 0 to (n – 1) • each value of l is called by a particular letter that designates the shape of the orbital – s orbitals are spherical – p orbitals are like two balloons tied at the knots – d orbitals are mainly like 4 balloons tied at the knot – f orbitals are mainly like 8 balloons tied at the knot l = 0, the s orbital • each principal energy state has 1 s orbital • lowest energy orbital in a principal energy state • spherical • number of nodes = (n – 1) l = 1, p orbitals • each principal energy state above n = 1 has 3 p orbitals – ml = -1, 0, +1 • each of the 3 orbitals point along a different axis – px, py, pz • 2nd lowest energy orbitals in a principal energy state • two-lobed • node at the nucleus, total of n nodes l = 2, d orbitals • each principal energy state above n = 2 has 5 d orbitals – ml = -2, -1, 0, +1, +2 • 4 of the 5 orbitals are aligned in a different plane – the fifth is aligned with the z axis, dz squared – dxy, dyz, dxz, dx squared – y squared rd • 3 lowest energy orbitals in a principal energy state • mainly 4-lobed – one is two-lobed with a toroid • planar nodes – higher principal levels also have spherical nodes l = 3, f orbitals • each principal energy state above n = 3 has 7 d orbitals – ml = -3, -2, -1, 0, +1, +2, +3 • 4th lowest energy orbitals in a principal energy state • mainly 8-lobed – some 2-lobed with a toroid • planar nodes – higher principal levels also have spherical nodes
