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Unit 2 – Symbols say WHAT?! ATOMIC THEORY AND NUCLEAR CHEMISTRY Atoms Atom- smallest particle of an element that retains the properties of that element. Atoms are electrically neutral A copper coin (penny) contains about 2.4 x 1022 atoms If you could line up 100,000,000 copper atoms side-by-side it would only equal 1 cm The radius of an atom is between 5.0 x 10-11m to 2.0 x 10-10m Properties of Subatomic Particles Protons, Electrons, and Neutrons are subatomic particles. Protons- a positively charged subatomic particle that is found in the nucleus of an atom- this has a charge of 1+ Electrons- a negatively charged subatomic particle that is found in the space outside the nucleus- this has a charge of 1- Neutrons- a neutral subatomic particle that is found in the nucleus of an atom Atomic Number Atomic Number- the number of protons in an element Ex) Oxygen- 8 , Hydrogen- 1 Atoms of different elements have different numbers of protons The atomic number identifies an element Since these are positive, they have to be balanced with negative charges. Therefore, the number of electrons equals the number of protons. Atomic Charge The charge of an atom is the difference in protons and neutrons Charge = p+ - e- Remember in neutral atoms these are the same Atoms with non-neutral charges are referred to as ions Mass Number Most of the mass of an atom is located in the nucleus Mass Number- the sum of the number of protons and neutrons in the nucleus of an atom Ex) A helium atom has two protons and two neutrons, so it has a mass number of 4 If you know the mass number and atomic number of an element, you can determine how many neutrons are in that atom Mass Number= protons(atomic number) + neutrons Carbon-14 Atomic mass = 14 Atomic number = 6 6 protons 8 Neutrons Tritium (hydrogen-3) Atomic mass = 3 1 proton 2 Neutrons Atomic number = 1 For each of the elements below, what is the atomic number and the number of neutrons? Their number is their mass number.(Use your reference tables) Carbon -14 Atomic Number = 6 Nitrogen – 13 Atomic Number = 7 Neutrons = 6 Fluorine-18 Atomic Number = 9 Neutrons = 8 Neutrons = 9 Rubidium -82 Atomic Number = 37 Neutrons = 45 Writing the Composition of Atoms Atomic number and mass number are written to the left of the element symbol AZX Write the symbol of the element (X) A= mass number Z= atomic number Gold with a mass number of 197 is written as 19779Au This can also be written as the element name- mass number Gold- 197 Isotopes Every atom of an element has the same number of protons, but not necessarily the same number of neutrons Isotopes- atoms of the same element that have different numbers of neutrons and different mass numbers These have the same atomic number, but different mass numbers because they have different number of neutrons Ex) oxygen has 8 protons and 9 neutrons- Oxygen has 8 protons and 8 neutrons- Atomic Mass It is difficult to determine the actual mass of individual atoms due to their small size and mass, so they compare the mass of an atom to a standard They compare it to a carbon-12 atom Atomic mass unit (amu)- 1/12 the mass of a carbon-12 atom In isotopes, this would be represented as the mass number The mass number is different for different isotopes, so they have different masses Average Atomic Mass The weighted average mass of the atoms in a naturally occurring sample of the element This is an average of the isotope masses according to abundance in nature You need to know the number of stable isotopes of an element, the mass of each isotope, and the % abundance of the isotope To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products Ex) Element X has two natural isotopes. The isotope with a mass of 10.012 amu (10X) has a relative abundance of 19.91%. The Isotope with a mass of 11.009 amu (11X) has a relative abundance of 80.09%. Calculate the atomic mass of this element. Find your known and unknown values X-10 10.012 X-11 11.009 amu x 0.8009= 8.817 amu Mass of element X 1.993 amu x 0.1991= 1.993 amu amu + 8.817 amu= 10.810 amu This should be closest to the isotope that is more abundant Ex) The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass=62.93 amu, and 30.8% for mass=64.93 amu. Calculate the average atomic mass of copper. Copper-63 62.93 Copper- 65 64.93 amu x 0.692= 43.54756 x 0.308= 19.99844 Mass of Copper 43.54756 + 19.99844= 63.546 amu On your own complete the following: Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%). Take a few minutes to calculate an answer. Answer: 79.9062 amu (78.92 x .5069) + (80.92 x .4931) Chlorine mass number exact weight Silicon percent abundance exact mass percent number weight abundance 35 75.77 28 92.23 37 24.23 29 4.67 30 3.10 The answer for chlorine: 35.453 amu The answer for silicon: 28.086 amu Nuclear Decay Radioactivity- the process in which an unstable atomic nucleus emits charged particles and energy The penetrating rays and particles emitted by a radioactive source Radioisotope/ radioactive isotope- any atom containing an unstable nucleus These isotopes are the source of any radioactivity Some common radioisotopes include: Uranium-238 Carbon-14 These isotopes can spontaneously change into other isotopes over time through nuclear decay During nuclear decay, atoms of one element can change into atoms of a different element altogether Ex) uranium-238 changes into thorium-234 This is done by losing or sharing electrons (emitting energy) Effects of Nuclear Radiation Without knowing it, you are exposed to nuclear radiation on a daily basis Background radiation- radiation that occurs naturally in the environment This comes from radioisotopes in Earth materials, trace amounts of radioactive elements, and cosmic rays Cosmic rays are streams of charged particles from outer space These sound dangerous, but levels are low enough to be safe Fission Fission- the splitting of an atomic nucleus into two smaller parts In nuclear fission, tremendous amounts of energy can be produced from very small amounts of mass Fission produces neutrons that can react with other atoms creating a chain reaction Converting Mass into Energy 30 years before fission was discovered, physicist Albert Einstein introduced the mass-energy equation E=mc2 Once fission begins, neutrons are released during each reaction creating a chain reaction In the chain reaction, neutrons released during the splitting of an initial nucleus trigger a series of nuclear fissions Modern Day Fission Fission can now be controlled and the energy can be released more slowly This is used in nuclear power plants to produce useful energy This reaction creates heat that is transferred to water creating steam This steam rotates a turbine that generates electricity Fusion Fusion- the process by which the nuclei of two atoms combine to form a larger nucleus This produces more energy than fission The sun and stars are powered by the fusion of hydrogen into helium These reactions require very high temperatures where matter exists as plasma Plasma- a state of matter in which atoms have been stripped of their electrons Scientists have not been able to make a fusion reactor due to the high temperatures that would be needed Types of Nuclear Radiation Nuclear radiation is charged particles and energy that are emitted from the nuclei of radioisotopes Common types of nuclear radiation include: Alpha particles Beta particles Gamma rays Alpha Decay Alpha particle- a positively charged particle made up of two protons and two neutrons (this is a helium nucleus) Symbol: 42He Alpha is also represented by the Greek letter α Nuclear reactions are expressed similar to chemical reactions Ex) Alpha decay of Uranium-238 Alpha particles are the least penetrating type of nuclear radiation. These can travel a few centimeters in air and can be stopped by a sheet of paper or clothing Beta Decay Beta particle- an electron emitted by an unstable nucleus 0-1e or β in a reaction This has an atomic number of -1 and mass number of 0 Due to their smaller mass and faster speed, beta particles are more penetrating than alpha particles- they can pass through paper, but they are stopped by metal Ex) Beta decay of Thorium-234 Gamma Decay Gamma ray- a penetrating ray of energy emitted by an unstable nucleus This has no mass and no charge These are similar to x-rays and visible light that move at the speed of light These are much more penetrating, it would take several centimeters of lead or meters of concrete to stop gamma radiation. The atomic and mass number stay the same but energy is released This often accompanies alpha or beta decay Ex) 23490Th → 23491Pa + 0-1e + γ Balancing Nuclear Equations Write a balanced nuclear equation for the alpha decay of polonium-210 Write a balanced nuclear equation for the beta decay of carbon14. Find the variables and solve. When nuclear radiation exceeds background levels, they damage the cells and tissues of your body The nuclear radiation ionizes atoms The bonds holding proteins and DNA together break causing the cell to function improperly Alpha, beta, and gamma are all types of ionizing radiation Alpha particles can cause burns, but tend to not do harm unless they are inhaled. Beta particles can penetrate tissues to do damage Gamma rays penetrate deeply exposing all organs to damage Half- Life Half-life- the time it takes for one half of a sample of radioisotope to decay Half-lives can vary from a fraction of a second to billions of years Here, half of the atoms have decayed and half stayed the same Some common half-lives are in a table on p. 299 Unlike chemical reaction rates, which vary with the conditions of a reaction, nuclear decay rates are constant Calculating Half-life 1st step- Calculate how many half-lives have passed. 2nd step- With the half-lives determine what fraction the sample has been reduced to. 3rd step- multiply the fraction by the mass of our original sample to determine how much of our original isotope is left 4th step- subtract the mass from step 3 from our original mass to determine the mass of our new isotope Remainder= mass x (1/2)n You have one-gram of iridium-182, which undergoes beta decay to form osmium-182. The half life of iridium-182 is 15 minutes. What will we have in our sample at the end of 45 minutes? Carbon-14 emits beta radiation and decays with a half-life of 5730 years. Assume you start with a mass of 2.00 x 10-12 g of carbon-14. How long is three half-lives? How many grams of the isotope remain at the end of three half-lives? Manganese-56 is a beta emitter with a half-life of 2.6 hours. What is the mass of manganese-56 in a 1.0 mg sample of the isotope at the end of 10.4 hours? Radioactive Dating In radiocarbon dating, the age of an object is determined by comparing the object’s carbon14 levels with carbon-14 levels in the atmosphere This gives an idea of how old a sample may be Nuclear Reactions in the Laboratory Transmutation- the conversion of atoms of one element to atoms of another These can be natural or artificial Scientists can perform artificial transmutations by bombarding atomic nuclei with high energy particles such as protons, neutrons, or alpha particles Rutherford was one of the first scientists to perform this experiment turning nitrogen into oxygen Transuranium Elements These are elements with atomic numbers greater than 92 (Uranium) All are radioactive and not found in nature Scientists synthesize a transuranium element by the artificial transmutation of a lighter element Neptunium was the first transuranium element synthesized from Uranium Most of these are synthesized purely for research, but a few are for commercial use Americium- 241 is created for smoke detectors Plutonium-238 decays powering a space probe Bohr’s Model of the Atom Bohr adapted Rutherford’s model to include discoveries of how the energy in an atom changes when it absorbs or emits light He proposed that an electron is found only in circular paths, or orbits, around the nucleus Energy level- the fixed energies an electron can have (each orbit) These are similar to rungs on a ladder- the lowest has the lowest energy and you can go up or down levels by gaining/losing energy It cannot be between energy levels Quantum- the amount of energy required to move an electron from one energy level to another energy level Energy is quantized This is not the same for each energy level The closer to the nucleus the electron is, the more energy it needs to move up an energy level The further away, the levels are closer together, so not as much energy is needed An electron can fall down a level releasing light energy in the form of a photon Bohr only created this model for hydrogen Light Study of light led to development of quantum mechanical theory and Schrodinger’s equation. Isaac Newton (1700’s) - light consisted of particles 1900 – light consisted of waves Parts of a Wave Amplitude- the wave height Wavelength- the distance between identical parts of a wave Frequency- number of wave cycles to pass a point in a certain amount of time Measured in hertz (Hz) Electromagnetic Radiation Radio waves, microwaves, infrared waves, visible light, ultraviolet waves, x-rays, and gamma rays This travels at 3.0 x 108 m/s (c) This radiation is a range of frequencies and wavelength in the spectrum c= λν (speed of light= wavelength x frequency) Frequency and wavelength are inversely proportional When sunlight goes through a prism, the different frequencies split into a spectrum of colors (rainbow) due to longer/shorter wavelengths Atomic Spectra Passing an electric current through a gas in a tube energizes the electrons in an atom and causes them to emit light When atoms absorb energy, electrons move into higher energy levels and they lose energy by emitting light when they return to a lower energy level Normal light contains all the wavelengths of light However, atoms only emit certain wavelengths of light, so when it splits through a prism it only shows certain color bands of light These lines are the atomic emission spectrum of the element No two elements have the same spectrum Electron Cloud Model Scientist have disproved Bohr’s theory that the electrons orbit like planets Electron cloud- a visual model of the most likely locations for electrons in an atom They move in an unpredictable manner The cloud is most dense where there may be electrons Scientists use the electron cloud model to describe the possible locations of electrons around the nucleus Quantum Mechanical Model This model determines the allowed energies an electron can have and how likely it is to find the electron is various locations around the nucleus This is based on probability, but does not limit electrons to energy levels This was developed by Erwin Schrodinger Atomic Orbitals Atomic orbital- a region of space in which there is a high probability of finding an electron These are found from using Schrodinger’s mathematical equations These come in different sizes and shapes S- orbitals are spheres P- orbitals look like dumbells (small in center and more balanced on each end) D- orbitals are similar to cloverleafs F- orbitals are much more complicated, because they have lobes on multiple planes Bohr’s energy levels are principle energy levels These are split into the orbitals Electron Configurations These show the arrangement of electrons around the nucleus of an atom Pxn P = principal energy level x = sublevel (s, p, d, f) n = # of electrons in whole sublevel Energy Levels Principal energy levels are divided into sublevels (s, p, d, f) Energy level 1 has 1 sublevel Energy level 2 has 2 types of sublevels Energy level 3 has 3 types of sublevels The number of the principal energy level equals the number of sublevels in that energy level Orbitals Sublevels are further divided into orbitals s sublevel: 1 orbital p sublevel: 3 orbitals d sublevel: 5 orbitals f sublevel: 7 orbitals Each orbital can have up to 2 electrons so s = 2 e-, p = 6e- , d = 10e- and f = 14e Writing Electron Configurations Rules for writing: 1. only two electrons per orbital 2. electrons enter the lowest energy orbitals first Can do electron configuration in two ways 1. diagonal diagram 2. from the periodic table Electron Configurations How to determine electron configuration with diagonals 1. always start with 1s, then 2s, 2p, 3s etc. 2. Figure out the number of electrons = atomic # if it is a stable atom 3. Look at your chart to see maximum # of electrons s = 2, p = 6, d = 10, f = 14 4. Write your electron configuration Diagonal Electron Configurations Examples: Sodium Manganese Bromine Exceptions to the Rule!- Honors Only Atoms are most stable when they have full or half full orbitals Ex) 3d5 is more stable than 3d3 This happens for many transition metals in the d block especially The energies of these two levels are very close together, so these elements will often not fill the s block in order to have a half full or full d block. Iron Silver Configurations from the Periodic Table The periodic table is broken up into blocks associated with sublevels (s, p, d, f) based on the location of the outermost electrons for each atom For example, any element in the first two columns have their outermost electrons in an s sublevel. So, these two columns are called the s-block - Li, Na, Mg, K, Sr, etc. Principal energy level is based on row number: For p and s blocks principal energy level = row number For d block: principal energy level = row# - 1 For f block: principal energy level = row# - 2 Examples Silver Uranium Xenon Noble Gas Configurations (abbreviated) Start with the noble gas before the element and write the configuration for that row Examples: Cobalt Tellurium Radium Electron Configurations of Ions Remember that ions are atoms with an unequal number of protons and electrons This gives them a non-zero charge Charge= protons-electrons A positive charge indicates the atom lost electrons A negative charge indicates the atom gained electrons To do a configuration for an ion, you add or subtract the charge from the total number of electrons Ion examples Mg+2 O-2 Cl-1 Orbital Diagram Rules Aufbau Principle e- enter orbitals of lowest energy first Orbitals of a sublevel are equal energy (s sublevel is lowest) Pauli Exclusion Principle an atomic orbital may describe at most 2 e- 1 or 2 e- may occupy an orbital and must have opposite spins clockwise or counter clockwise. Hund’s Rule Represented by arrows ↑↓ for orbitals of equal energy, one e- enters each orbital until all orbitals contain 1 e- before pairing Stable configurations have half- full or full orbitals Examples Hydrogen Krypton Titanium Atomic Radius/Size Atomic Radius- one half the distance between the nuclei of two atoms of the same element when the atoms are joined Atomic size increases from top to bottom in a group and decreases from left to right across a period Going down a group, the number of energy levels increases. The electrons in the outermost level are shielded by electrons closer to the nucleus from the pull of the protons.- this is what causes the size to increase down a period Across a group, the number of energy levels does not change, so the shielding effect is not increased. You increase the number of protons, so this pulls the outer electrons closer to the nucleus.- this causes the size to decrease across a period. Ions & Ionic Radius Ion- a charged particle that is produced after an atom gains or loses an electron Cation (+)- these ions lose electrons to become positively charged These tend to be metals Since an electron was lost, there are more protons than electrons. The electrons are pulled closer to the nucleus making the ion smaller than the element. Anion (-)- these ions gain electrons to become negatively charged These tend to be nonmetals Since an electron was gained, there are more electrons than protons. The electrons cannot be pulled as close to the nucleus making the ion larger than the element.
