Unit 1 Lecture 2
*Note: The charges are being compared
Neutron relative to0each other, they do have
specific values (C/g), see Thomson slides
• A neutral atom has an equal amount of protons and electrons • Protons and neutrons have equal masses (they are actually a little bit different, but this is several decimal places out and using a very expensive measuring device
• The mass of an electron is MUCH less than the mass of a proton or neutron, thus the mass of electrons is rarely taken into account, and we say it has a relative mass of 0.
Mass and Atomic
• Mass Number (A) o Equal to the number of protons and neutrons o Can have the same element, but with different mass numbers…these are ISOTOPES o NOT on the periodic table box
• Atomic Number (Z) o The number of protons in an atom o In a neutral atom this is also the number of electrons o Top number in periodic table box Check yourself! 1. What is the mass of this atom of carbon?
2. How many electrons does a neutral carbon atom have?
1. 14 amu
2. 6 electrons
(b/c all carbon have 6 protons and to make it neutral you would need 6 negative electrons)
3. How many neutrons are in this atom of carbon? 3. 8 neutrons
Mass # is 14, 6 of which comes from the protons…14-6=8
4. What does 12.011 represent? 4. Atomic Mass (aka…average atomic mass….more on this soon!) More practice! 1. What would the full name of this isotope of hydrogen be called?
2. How many electrons?
3. How many neutrons?
4. How much would 2 atoms of hydrogen weigh? 1. Hydrogen-1
Use full name followed by the mass number…also more on isotopes in a minute!
2. 1 electron
Assume it is neutral unless otherwise stated
3. 0 neutrons
• Isotopes of an element have the same number of protons, but different numbers of neutrons
• Isotopes of an element exhibit identical chemical behavior o For example, both hydrogen-1 and hydrogen-2 can bond with oxygen to form water o FUN FACT! Hydrogen-2 (deutrium) is actually the hydrogen found in heavy water…it’s called heavy b/c with a large enough sample, it will increase the overall mass of water!
A Note on Atomic
• The mass of one carbon-12 atom is exactly
• All other atoms are measured on a scale relative to carbon-12 o You will notice that the nucleons (protons and neutrons) in all other atoms are not exactly 1 amu, but pretty darn close! o We will learn in nuclear chemistry later that mass is converted into energy when a nucleus binds together according to the formula E=mc2
…strong nuclear force, etc
Average Atomic Mass
• The atomic mass given to you on the PT (periodic table) is an average of all the isotopes of that element • Given a large enough sample, the average mass of atoms of a given element is always the same for that given element. o The ratio of isotopes for each given element is constant
Average Atomic Mass cont’d • Notice if you took a normal average of all of the isotopes of carbon, the atomic mass would be 13 amu. o However, on the PT you see that carbon’s atomic mass is 12.011amu
• That is because we take a weighted average o Each isotope of carbon is not found in equal amounts in nature
Relative abundance of each type of carbon in nature
• Used to compare the masses of isotopes
• Atoms from a pure samples of an element are ionized
(some electrons removed making them ions) and accelerated through a magnetic field
• Isotopes with smaller masses experience a greater degree of deflection o Positive ions go towards the negative plate (acceleration= Force/ mass) o The