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Elements and Periodic Trends Chemistry Mr. Treanor Atomic Theories Democritus – first known theory of atoms: “Matter is composed of very small solid particles that are indivisible.” Dalton’s Theory – five points in his theory explain: – the law of conservation of mass – the law of definite proportions – the law of and multiple proportions More Atomic Theories Experiments with cathode-ray tubes showed that atoms were composed of even smaller particles. Millikan discovered the mass and charge of the electron. Rutherford’s gold foil experiment led to the discovery of the nucleus. These all led to the Bohr Model. Quantum Model Electrons exist in orbitals, not orbits. Orbitals are regions around the nucleus where there is a 90% probability (as they are usually shown) that electrons will be found. Heisenberg Uncertainty Principle: it is impossible to know the speed and location of an electron at the same time Schrödinger Wave Equation: mathematical equations for the shapes of orbitals Shapes of orbitals l = 0: these are s orbitals (spherical) l = 1: these are p orbitals (pair) l = 2: these are d orbitals (drops/donut) l = 3: these are f orbitals (funky) The f orbitals (funky shapes) Electron Configurations Aufbau principle: electrons go into the lowest energy orbital that has room Pauli exclusion principle: no two electrons in the same atom have the same set of quantum numbers Hund’s rule: electrons fill up empty orbitals first More about configurations The maximum number of electrons in each energy level is found using the quantum number n – main energy level The number of electrons can be up to 2n2 This gives 2,3,18, and 32 for n= 1,2,3, and 4 The highest occupied level is the main energy level with the highest n Inner-shell electrons are the ones that are not in the highest occupied level Electron Configuration Notation Used to show the locations of all electrons This is the order electrons fill the orbitals: 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d,7p The small numbers are the number of electrons in that particular orbital Example: Na is 1s22s22p63s1 – There are 2 electrons in the 1s orbital, 2 in the 2s orbital, 6 in the 2p orbital, and 1 in the 3s orbital The highest occupied level is 3 The inner shell electrons are in 1s,2s, and 2p How to tell which element? If the outer electrons are only s, it is an sblock element. One s electron means group 1; 2 is 2. If the outer electrons have p orbitals, it is a p-block element. 1 p electron is group 13, 2 is 14, and so on until the p orbitals are full with 6. These are group 18. The entire energy level is full. That is why noble gases are unreactive. If there are s and d, it is a d-block metal. Brief History Mendeleev invented the periodic table and arranged it by atomic mass and group properties. Moseley, working in Rutherford’s lab and playing with x-rays, arranged the periodic table by atomic number. They did not know about the noble gases yet. A look at the blocks •s-block – metals that easily lose electrons •d-block – metals that lose different numbers of electrons depending on the other element •p-block – mostly nonmetals that gain electrons Periodic Group Properties Group 1 Alkali metals, most reactive metals, only 1 outer s electron Group 2 Alkaline earth metals, 2 outer s electrons (no other outer electrons) Transition metals, the ones we call metals (gold, copper, iron, silver) Mostly nonmetals & semiconductors Halogens, most reactive nonmetals, 5 outer p electrons Noble gases, all main energy levels full D-block P-block Group 17 Group 18 Atomic Radius Gets bigger as you go down a group (more shells) Gets smaller as you go across a period (more protons with same number of shells) Measured as half the distance between 2 atoms in a pure sample of a substance Ionization Energy (IE or IP) The energy required to remove an electron from a neutral atom is the 1st ionization energy, or IE1. The energy required to remove a second electron after the first ionization is the 2nd ionization energy, or IE2. There is also an IE3, IE4, and so on. Electron Affinity (EA) The ability of an atom to gain electrons Negative numbers mean more energy Electronegativity It is defined as the relative ability of an atom in a molecule to attract electrons towards itself. The atom with the higher electronegativity will dominate the electrons. Electronegativity (EN) Values What causes Electronegativity? Protons in the nucleus are attracted to valence electrons in OTHER ATOMS If there are a lot of electrons in the way, the electronegativity is lower. This happens when there are more energy levels (shells) If there are more protons, they will be a stronger pull and higher electronegativity. This happens when the atomic number is higher. Atomic Radii 1. Which atoms have less shells? lower 2. Of those, which ones have more protons? To the right 3. Explain why this agrees with the EN values. The elements with less shells and more protons have the highest electronegativity values. The Tug-O-War How electronegativity works: The atom with the highest EN is the strongest and wins. Valence Electrons Outer s and p electrons For group 1, it is 1; for group 2 it is 2. For groups 13 – 18, it is the group number -10