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Atoms: The Building Blocks of Matter 3 Guiding Principles Chemistry: 1. Law of Conservation of Mass/Matter 2. The Law of Definite Proportions 3. The Law of Multiple Proportions Dalton’s Atomic Theory Problems with Dalton’s Atomic Theory: Matter can neither be created nor destroyed o You can not make an atom appear or disappear o The atoms you start with in a reaction must be accounted for in the product - A chemical compound contains the same elements in the same ratio regardless of mass of sample o A tsp of water, a glass of water, and a swimming pool of water all have the ratio of 2H : 1 O regardless of size of sample - If two or more different compounds are composed of the same elements the ratio of the elements in those compounds will always be simple whole number ratios o You can never have half an atom or part of an atom the ratio of elements must always be the lowest whole numbers. 1. All elements are made of tiny indivisible atoms. Atoms can not be created nor destroyed 2. Atoms of the same elements are identical in size, mass, and other properties. Every element has its own unique atom. 3. Atoms can combine to form compounds 4. Atoms can combine, separate, or be rearranged during a chemical reaction - 1. We have discovered that atoms can be broken down into subatomic particles: protons, neutrons, and electrons - We have also learned that quantum particles exist: quarks, leptons, positrons etc. 2. Atoms of an element are not always identical. Some elements have atoms with different masses. These atoms are called isotopes The Subatomic Particles: The Electron - Discovered in 1897 by J.J. Thomson - Ran a charge through the cathode-ray tube and a beam of light appeared. Checked to see if particles in the beam were positive or negative Knew that like charges repel, opposite charges attract When he placed a positively charged plate near the tube, the beam was attracted to it. When he placed a negatively charged plate near the tube, the beam was deflected away. Cathode Ray Tube Experiment - o Based on the above he concluded the beam contained negatively charged particles o Called the negative charges Electrons Facts about electrons ---------------------------------------THE PROTON The Rutherford Gold-Foil Experiment - Electrons are small negatively charged particles Electrons are 10,000 times smaller than protons/neutrons (9.109 x 10 -31 Kg = 0 amu) - Electrons are responsible for the bonding that occurs among atoms - Electrons float around the nucleus in an area called the electron cloud - Electrons carry a full 1- charge -------------------------------------------------------------------------------- Discovered in 1911 by Ernest Rutherford - 3 important findings of Gold-Foil Experiment Facts about the Proton Major contribution of the proton The atomic number Wanted to know if and where any positive particles existed in the atom - Placed radioactive Polonium in a lead box. o Polonium gives off tiny positive particles called alpha particles - If positive alpha particles escape the box and pass straight through the gold-foil to the X-ray film, the atom must not contain any positive charges (like charges repel) - If some of the positive alpha particles are deflected backward, then the atom must contain positive charges – this theory turned out to be true o Most particles went straight through , but a few were deflected ------------------------------------------------------------------1. Atoms contain positively charged particles – Protons 2. Atoms are mostly empty space in which electrons can roam about 3. Atoms are located in a densely-packed, positively charged, central core – called the Nucleus. - The proton is a positively charged particle- carries a full 1+ charge - The proton lives in the nucleus of the atom - The proton has a mass of 1.673 x 10-27 kg = 1amu o Amu stands for: atomic mass unit o 1 amu = 1.67 x 10 -27 kg - A proton and an electron have opposite charges which can cancel each other out and make the overall charge neutral -----------------------------------------------------------The number of protons in an atom are unique to each element - the number of protons in the atom o serves as the atoms fingerprint -------------------------------------------------The Neutron o identifies one element from another o elements are organized on the periodic table according to their atomic number o atomic # = # of protons ------------------------------------------------------------------------------------- discovered by James Chadwick in 1932 - Has approximately the same mass as a proton o 1.675 x 10 -27 kg = 1 amu - Found in the nucleus - Adds weight/heft to the nucleus - Has no charge (composed of equal amounts of positive and negative quantum particles) = neutral - Partially responsible for radiation - The number of neutrons in an element may be different creates isotopes The Nucleus: The mass of atom is located in the nucleus The nucleus has a net positive charge Contains protons and neutrons Held together by atomic forces Densely- packed core: contains all the mass of the atom, but takes up very little space (if atom is the size of a football stadium, the nucleus would be a marble sitting in the middle of the field) 2 Masses associated with an element - Mass number - Average atomic mass - How is the mass of an atom determined? Mass Number - Formula - Solving for neutrons or protons: o Average Atomic Mass Let’s review what isotopes are. - Determined by the total number of protons and neutrons in an atom Mass number = # protons + # neutrons Explanation: I know that… o 1 proton weighs 1 amu o 1 neutron weighs 1 amu o 1 electron has no real mass (negligible) so we do not count it o If I have 4 protons they must weigh 4 amus and if I have 6 neutrons they must weigh 6 amus. Therefore the mass of the atom is the total of 4 amu + 6 amu = 10 amu If I know the mass of an atom is 39 amus and I know that its atomic number is 19, I can solve for the number of neutrons by plugging the given values into the formula. Mass # = #protons + # neutrons 39 = 19 + X Solve for X X = 39 - 19 = 20. There are 20 neutrons in this atom. Isotopes: Atoms of the same element that have different masses. o Isotopes of the same element must always have the same number of protons o Isotopes can have different numbers of neutrons and different mass numbers (30 represents the mass of isotope E, the 15 represents isotope E’s atomic number 3017E 3515E 3512E Example 1: 3015E A. B. C. D. o Example of an Isotope Question Which two letters represent isotopes of the same element? Answer: -----------------------------------------------Let’s review the common math formula for finding the average of a list of numbers: 3015E The correct answer is A and C. A and C both have an atomic number 15, which signifies that they are the same element. However, both A and C have different masses they are isotopes of that element. ------------------------------------------------------------------------------------------------Most of us learned that to find the average of a given series of numbers, we must add the numbers together and divide by the total number of numbers. For Example: Find the average of 3, 5, 4, and 2 3 + 5 + 4 + 2 = 14 divide by the number of numbers (4) 14/ 4 = 3.5 therefore 3.5 is the average But why do we divide by 4? We divide by 4 because each number in our series makes up an equal percentage in our average. In other words each number makes up 25% of the total. Let’s find the average in a different way. This time instead of dividing by 4 at the end, let’s multiply by the percentage weight of each number. (4 x 25/100) + (3 X 25/100) + (5 X 25/100) + (2 x 25/100) = 3.5 The number is the same. What do you do if the numbers being added together are not represented equally? Formula for average atomic mass: Example 1: What is the average atomic mass of carbon if C-12 makes up ~ 98% of all carbon atoms and C-14 makes up only ~2%? In nature, the majority of the isotopes of an element do not exist in equal amounts. For example Carbon -12 makes up ~ 98% of all carbon atoms, while Carbon -14 only makes up ~2% of the total carbon atoms. If I were to find the average of both isotopes in the traditional way, the average mass would be 13. This, however, is not correct; the majority of carbon atoms have a mass of 12, so why should Carbon-14 weigh so heavily on the average. To deal with this inequality we use the average atomic mass formula. (mass of isotope A x % abundance) + (mass of isotope B x % abundance) + etc. 100 100 Solving Example 1 (12 amu x 98) + (14 amu x 2 ) = 11.76 + .28 = 12.04 100 100