Unit 2: Symbols Say WHAT?

The Three Subatomic Particles After various molecular models of the atoms had been tested, it was determined that three subatomic particles made up the atom Protons Neutrons Electrons

Protons Protons are found in the nucleus Protons have an actual charge of +1.6 x 10-19 C and a relative charge of +1 The actual mass of a proton is 1.67 x 10-24 g The relative mass of a proton is 1 atomic mass unit (amu) The symbol is p+

Neutrons Neutrons are found in the nucleus Neutrons have an actual charge of 0 C and a relative charge of 0 The actual mass of a neutron is 1.67 x 10-24 g The relative mass of a neutron is 1 atomic mass unit (amu) The symbol is n0

Electrons Electrons are found outside the nucleus Electrons have an actual charge of   -1.6 x 10-19 C and a relative charge of -1  The actual mass of an electron is   9.11 x 10-28 g The relative mass of a electron is 1/2000 atomic mass unit (amu) The symbol is e-

Atomic Number Each element has a certain number of protons in its nucleus The number of protons in the nucleus is called the atomic number Each element has its own atomic number because each element has its own, unique number of protons Note: That an element only changes if the atomic number changes. Which means that if the number of protons changes it changes elements

On Your Own Which element has a) 87 protons b) 35 protons c) 50 protons d) 92 protons e) 8 protons f) 19 protons

On Your Own: Determine the number of protons in the following atoms as well as each atom’s identity a) 6 electrons b) 14 electrons c) 72 electrons d) 55 electrons

Mass Number = Protons + Neutrons Mass Number: the number of protons and neutrons in an atom added together Mathematically Mass Number = Protons + Neutrons

Example What is the mass number of an atom with 16 protons and 16 neutrons?

Determine the Mass Number for the following atoms On Your Own Determine the Mass Number for the following atoms a) 17 protons and 18 neutrons b) 11 protons and 12 neutrons c) 1 proton and NO neutrons d) 3 protons and 4 neutrons

Nuclide Symbols Mass Numbers (This is represent by an A) are written in the upper left preceding the chemical symbol Atomic Number (This is symbolized by a letter Z) is written directly under the mass number

Example: Write the correct nuclide symbol for an element with 51 p+ and 71 n0

Practice Draw the following Nuclide Symbols for the following: Sr Os Zn I Cs K

Practice Write the following elements in proper Nuclide Symbols: Pt Fr At K

Warm-Up 2/10/2015 -Write the correct Nuclide Symbol for the following elements: *19 p+, 20 n *82 e-, 125n *Mass # 238, neutrons= 146

Isotopes Atoms of the same elements have the same number of protons HOWEVER there may be different numbers of neutrons and different mass numbers When an element’s atom has different numbers of neutrons, it is said to have isotopes

Hydrogen’s Three Isotopes Hydrogen has the following isotopes: Protium-a hydrogen atom with one proton and NO neutrons Deuterium-a hydrogen atom with one proton and only one neutron Tritium-a hydrogen atom with one proton and two neutrons

The atomic mass unit is based on the relation of standard carbon-12 Average Atomic Mass Atomic mass is the mass of an atom expressed in atomic mass units or amu The atomic mass unit is based on the relation of standard carbon-12 Begin 2nd here on 2/13/2012

Average Atomic Mass Continued Carbon-12 has a mass of 12.000 00 amu Example: If an atoms weighs half as much as carbon-12, its atomic mass will be 6.000 amu Begin 4th 10-13-2000 begin 5th, begin 7th

Example: If an atom weighs four times as much as carbon-12, it will have a mass of 48.000 00 amu

What is Atomic Mass? The atomic mass that is reported in the periodic table is a weighted average based on the relative abundance of each element

Relative abundance refers to how common the isotope occurs in nature Percent Abundance which refers to how many of each isotope are in every hundred

To Determine Avg. Atomic Mass 1) First convert relative abundance (%) to decimal equivalent 2) Multiply mass (in amu) by decimal equivalent 3) Add the numbers together 4) The sum (in amu) is the average atomic mass

For example, an element has two naturally occurring isotopes For example, an element has two naturally occurring isotopes. One isotope has a relative abundance of 19.91% and a mass of 10.012 amu. A second isotope has a relative abundance of 80.08% and a mass of 11.009 amu. Calculate the atomic mass

Example For example, an element has two naturally occurring isotopes. One isotope has a relative abundance of 92.58% and a mass of 7.02 amu. A second isotope has a relative abundance of 7.42% and a mass of 6.02 amu. Calculate the atomic mass

Additional Example Calculate the average atomic masses for the following: Isotope: Rel. Abund. Rel. Mass hydrogen-1 99.985% 1.008 hydrogen-2 0.015% 2.014 Begin 3rd 10-16-00, 4th 1.007 amu 1.0 amu (s.f.)

Practice Titanium has five common isotopes: If the abundance of Ti- 46 is 8.0%, Ti-47 is 7.8 %, Ti-48 is 73.4 %, Ti-49 is 5.5% and Ti- 50 is 5.3 %. What is the average atomic mass of titanium?

Isotope Rel. Abund. Actual Mass O-16 99.762 15.995 O-17 0.038 16.999 Determine Avg. Atomic Mass for oxygen: Isotope Rel. Abund. Actual Mass O-16 99.762 15.995 O-17 0.038 16.999 O-18 0.200 17.999

Review 2/11/2015 What is an Isotope? What are hydrogen’s three isotopes names? What is the difference between average atomic mass and mass number?

Review Practice Rubidium is a soft, silvery-white metal that has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?

Practice Begin 3rd 10-16-01 & 4th

Practice Element Atomic # Mass # p+ n0 e- Ag 110 19 40 127 53

Practice

Finding % Abundance Silver( Atomic Weight 107.868) has two naturally- occurring isotopes with weights of 106.91 and 108.90. What is the percentage abundance of the lighter isotope?

Practice Oxygen comes in two stable isotopes, Oxygen- 16 and Oxygen- 18. The molar mass of Oxygen- 16 is 15.99 amu. The molar mass of Oxygen-18 is 17.99 amu. Determine the percent abundance of each isotope.

Finding Percent Abundance The element indium exists naturally as two isotopes. 113In has a mass of 112.90 amu, and 115In has a mass of 114.90 amu. The average atomic mass of indium is 114.82 amu. Calculate the percent relative abundance of the two isotopes of indium. Candium Activity Afterwards

Review Antimony has two naturally occurring isotopes. The mass of antimony-121 is 120.90 amu and the mass of antimony-123 is 122.90 amu. Using the average mass from the periodic table, find the abundance of each isotope.

Practice Europium has two stable isotopes: Eu-151 with a mass of 150.92 amu and Eu-153 with a mass of 152.92. If elemental Europium is found to have an average mass of 151.964 amu on earth, calculate the percent of each of the two isotopes