8/29/2011. The Greek Philosophers. Atomic Structure & The Periodic Table. Dalton s Atomic Theory (1808) J. J. Thomson. Thomson s Experiment

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1 Atomic Structure & The Periodic Table The Greek Philosophers Democritus believed that all matter is made up of tiny particles that could not be divided Aristotle -- thought that matter was made of only four elements Fire Earth Air Water Dalton s Atomic Theory (1808) All elements are composed of atoms All atoms of the same element have the same mass, and atoms of different elements have different masses* Compounds contain atoms of more than one element In a particular compound, atoms of different elements always combine in the same way J. J. Thomson Provided the first evidence that atoms are made up of even smaller particles. Plum Pudding model electrons were evenly spread out among a positively charged mass (chocolate chip ice cream) Thomson s Experiment Hypothesis The beam was a stream of charged particles that caused the air to glow. Experiment Positive and negative plates were put on either side of the tube. Results The beam was bent towards the positive plate. Conclusion There are negative particles being given off of the atoms that are attracted to the positive plate. (Atoms are made of smaller pieces!) 1

2 Ernest Rutherford s Gold Foil Experiment Hypothesis: that alpha particles would pass straight through a thin sheet of gold Results The alpha particles did not pass straight through, but instead were deflected in various directions. Conclusion The positive charge of an atom IS NOT evenly spread out, but concentrated in a very small area (nucleus). Subatomic Particles Proton (p + ) Positive charge Found in nucleus Neutron (n) Neutral (no charge) Found in nucleus Approximately the same mass as a p + Electron (e - ) Negative charge Found orbiting around the nucleus MUCH smaller than p+ or n (1/1836) Atomic Number (Z) The number of protons found as a part of the nucleus in an atom Unique for each element Cannot change without changing the type of element Mass Number The total number of particles that make up the nucleus Mass number = p + + n n = mass number p + Atoms of the same element may have varying mass numbers because the number of neutrons can vary 2

3 Isotopes Atoms of the same element with different mass numbers (numbers of neutrons) Ways to write isotopes Element mass number (Carbon 12) Symbol mass number (C 12) Mass number Symbol ( 12 C) Ways to write isotopes, cont. Wave Nature of Energy MassNumber Atomic Number Symbol ( 12 6 C) Electromagnetic (EM) Radiation any form of energy that radiates in all directions from a single source consists of oscillating electric and magnetic fields that carry energy EM Spectrum wave formula c = discovered by James Maxwell c = speed of light (true for ALL EM radiation) c = 3.0 x 10 8 m/s = frequency the number of waves passing a given point during a unit of time units are per seconds = s -1 = Hertz (Hz) = wavelength (m) 3

4 What is the wavelength of a radio wave broadcasted at MHz? (2.901 m) blackbody radiation heated solid objects emit visible light where intensity and color depend on temperature black because there is no light emitted before heating Wien-Planck law as the temperature of a black body increases decreases increases E increases cannot be explained by classical physics quantum (Max Planck ) a fixed amount of energy the smallest amount of energy that can be emitted or absorbed as EM radiation E = h h = Planck s constant = x Js E = energy (Joules, J) one quanta = h, two quanta = 2 h photon energy packet that behaves like a particle photoelectric effect clean metal surfaces exposed to light will emit electrons emitted electron = photoelectron 4

5 What is the energy of one photon of yellow light ( = 589 nm)? 3.37 x J/photon Einstein extended Planck s quantum theory to say that energy has mass! E = mc 2 E = energy of a photon (J) m = mass of photon (kg) c = speed of light Compton s Theory ties together Planck s and Einstein s work E = mc 2 = h What is the mass of a photon of violet light ( = 415 nm)? 4.79 x kg m = hv/c 2 (c = ) m = h/ c De Broglie (Louis) equation m = h/ v m = mass (kg) h = Planck s constant = wavelength (m) v = velocity (m/s) Calculate the wavelength of a 2.53 g bullet traveling 320. m/s x m If mass is large, wavelength is small Newtonian mechanics applies If mass is small, wavelength is large Quantum mechanics applies 5

6 Wave-Particle duality Planck and Maxwell work toward wave theory Einstein work toward particle theory EM has characteristics of both a wave and a particle at all times Niels Bohr 1 st postulate atom has only certain allowable (quantized) energy states depend upon energy level occupied by e - Energy levels identified with integers (n=1, 2 ) energy of an electron in energy level n E = -kz 2 /n 2 k = x J (Rydberg constant) Z = atomic number n = principal quantum number (energy level) 2 nd postulate atom doesn t radiate energy in one of it s energy states 3 rd postulate atom changes energy states by absorbing or emitting photons of specific frequencies if a photon is absorbed, an electron is promoted to a higher energy level (called an excited state) E > 0 Calculate the energy absorbed by an electron jumping from the 1 st energy level to the 5 th x J When an electron relaxes back to its ground state, a photon is emitted E < 0 E 2.179x n 2 i 1 n 2 f 6

7 Calculate the energy of an electron moving around a H atom in the 2 nd energy level x J Energy Levels The possible energy an electron can have Similar to steps Higher steps have more energy Going down a step means energy was released Energy Levels, cont. Seven possible energy levels that correspond to the rows on the periodic table Also called shells Bohr s Model of the Atom Each circle represents an energy level 1 st level: 2 electrons 2 nd level: 8 electrons 3 rd level: 18 electrons 4 th level: 32 electrons Erwin Schrödinger Developed mathematical model to describe the motion of electrons Work leads to electron cloud model Orbital Region of space where an electron is likely to be found Shapes are created by taking many pictures of the electron s location over a period of time 7

8 Pauli Exclusion Principle Only two electrons will occupy the same orbital A series of the same type of orbital is also called a sub-level or subshell. s-orbital Shaped like a sphere Only one type (b/c it can only point one direction!) Electron Probablity Map p orbital p - orbital Has one shape that can point three ways Total of 3 types of p-orbitals (p x, p y, p z ) Electron Probability Maps d orbitals d - orbital Five total types 8

9 f - orbitals Seven possible types How Orbitals Interact Electron Configurations (More sophisticated than Bohr s!) Know the total number of electrons in the atom! Use the diagonal rule Aufbau Principle Electrons fill orbitals that have the lowest energy first Follow the DIAGONAL RULE! The Diagonal Rule 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Exceptions to the Diagonal Rule Expected Found Cr [Ar]4s 2 3d 4 [Ar]4s 1 3d 5 Mo [Kr]5s 2 4d 4 [Ar]5s 1 4d 5 Cu [Ar]4s 2 3d 9 [Ar]4s 1 3d 10 Ag [Ar]5s 2 4d 9 [Ar]5s 1 4d 10 Au [Ar]6s 2 5d 9 [Ar]6s 1 5d 10 9

10 Orbital Diagrams Shows the placement of electrons in the orbitals Ex. Sodium 1s 2 2s 2 2p 6 3s 1 Hund s Rule Electrons will fill each orbital on a sublevel before pairing with each other. Ex. Carbon 1s 2s 2p 3s Diamagnetic all electrons are paired s 2, p 6, d 10, f 14 Paramagnetic One or more electrons is unpaired Quantum Numbers Principal Quantum Number (n) n = energy level (1, 2, 3, ) Azimuthal Quantum Number (l) l = type of orbital (s=0, p=1, d=2, f=3) Magnetic Quantum Number(which orbital) m = from - l to + l Electron Spin (first or second electron) s = ½ What are the quantum numbers for the 3 rd and 6 th electrons in neon? N 1s 2 2s 2 2p 6 Electron Configurations on the Periodic Table 1s 2s 2p 3rd electron n=2, l=0, m=0, s=+½ 6 th electron n=2, l=1, m=-1, s=-½ 10

11 Shorthand electron configurations Use Noble Gas shortcuts ONLY! Kr (36 e - ) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Isoelectronic series Species that have the same electron configuration E.g. O 2-, F -, Ne, Na +, Mg 2+ Mo (42 e - ) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 4 [Kr] 5s 2 4d 4 John Newlands Arranged the 16 known elements in order of increasing atomic mass Dmitri Mendeleev Arranged the periodic table by increasing atomic mass and stacking elements with similar properties in the same column Was able to predict the properties of elements not yet discovered by looking at the blank spaces in his table Henry Moseley Fine-tuned the periodic table by placing the elements in increasing atomic number (protons were not known during Mendeleev s time) 11

12 Organizing the Periodic Table Metals are left of the stair-step line Non-metals are to the right of the stair-step line Metalloids (have properties of both) are found adjacent to stair-step line Other Periodic Arrangements 1 Atomic Information from the Periodic Table Atomic Number H Atomic Mass Element Symbol Atomic Mass The weighted average of all possible isotopes for an element Atomic Mass = (%A)(Mass of A) + (%B)(Mass of B) Period A row in the periodic table corresponding to the energy level on which the electrons exist 12

13 Group A column on the periodic table in which the elements have similar properties Similar properties are created by similar electron configurations 1A A Groups = Representative Elements 2A 3A 4A 5A 6A 7A 8A B groups = transition elements Group 1A are the alkali metals (but NOT H) Group 2A are the alkaline earth metals H These are called the inner transition elements, and they belong here Group 8A are the noble gases Group 7A is called the halogens Valence Electrons Electrons that are in the outermost shell Involved in bonding with other atoms Same amount as A group # Examples: H has 1, N has 5, P has 5 13

14 Octet Rule Atoms are most stable when they have full outer shells 8 electrons (s 2, p 6 ) for most representative elements (s2 for H, He) Ion an atom (or group of atoms) that has gained or lost electrons anion negative charge gained electron(s) tend to be larger than parent atom cation positive charge lost electron(s) tend to be smaller than parent atom Monoatomic Ions Ions from single atoms Charge can be predicted by location on the periodic table Cation positively charged ion lost e Anion - negatively charged ion gained e - Atomic Size Atomic Size - Group trends As the atomic number increases H Li } Radius Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule. each atom has another energy level, so the atoms get bigger. Na K Rb 14

15 Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. Electrons are in the same energy level. But, there is more nuclear charge. Outermost electrons are pulled closer. Ionization Energy Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom). Removing one electron makes a 1+ ion. 2 nd IE > 1 st IE Na Mg Al Si P S Cl Ar IE for non-valence electrons >>> 1 st IE IE - Group trends As you go down a group, the first IE decreases (less energy required, easier to remove electron) because the electron is further away from the attraction of the nucleus Increased sheilding Shielding The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. Second electron has same shielding, if it is in the same period IE - Period trends All the atoms in the same period have the same energy level. But, increasing nuclear charge So IE generally increases from left to right (higher energy, more difficulty to remove electron) Trends in Electronegativity Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. They share the electron, but how equally do they share it? The higher the EN the stronger an atom is at attracting electrons. 15

16 Electronegativity Group Trend The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has. Thus, more willing to share. Low electronegativity. Electronegativity Period Trend Metals let their electrons go easily Thus, low electronegativity Nonmetals want more electrons. Try to take them away from others High electronegativity. Electron Affinity The amount of energy RELEASED when an electron is added to an atom the more negative, the more energy that is released, the more likely an atom will add an electron Trends in three atomic properties Trends in metallic behavior 16

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