Atoms – a closer look at elements As mentioned before, atoms are composed of protons, neutrons, and electrons

So, what does it mean to define an element? An element is defined as the number of protons that it has in its nucleus. Number of protons = Atomic Number The atomic number is the same for all atoms of the same element.

Atomic number Equal to protons Equal to electrons IF atom is neutral Determines an elements placement on the periodic table Represented by letter Z So Z = protons = electrons (…if neutral)

Practice on your own! How many protons does titanium (Ti) have? Antimony has 51 protons, what is its atomic number? What abbreviation is given to Antimony? Magnesium (Mg) is a neutral atom, how many electrons does it have?

Ag Mass number (A) Mass number = protons + neutrons To find neutrons given protons and mass number: Neutrons = mass number - protons 108 Ag This is called a nuclear symbol 47 Top is mass number Bottom atomic number Note: Both numbers on left

Atomic Mass Unit (amu) 1 atomic mass unit (amu) is defined as 1/12th the mass of a carbon-12 atom 1 amu is nearly (but not exactly) equal to 1 proton or 1 neutron

Practice Problem What is the mass number of an element with 10 protons and 10 neutrons? What element are we dealing with? How would you represent this symbolically? 20 amu Ne 20 10

More Practice How many protons, neutrons and electrons does the following neutral atoms have? Fe 26 protons, 30 neutrons, 26 electrons Xe 54 protons, 77 neutrons, 54 electrons 56 26 131 54

Two More – now with a charge 16 protons 16 neutrons 18 electrons 20 protons 21 neutrons 19 electrons

Look at the masses on the Periodic Table So, if the atomic mass is equal to the sum of protons and neutrons, why isn’t the mass a whole number on all Periodic Tables of the Elements? Even though an atom has to have a certain number of protons, the number of neutrons can vary slightly. These different versions of the same atom are called… Isotopes

Isotopes Same number of protons but different number of neutrons Examples: Carbon-13 vs. Carbon-12 13 & 12 represent the weight of the particular carbon atom. Isotopes with more neutrons will have greater mass. Isotopes of an atom have the same chemical behavior.

Average Atomic Masses Relative atomic mass: average masses of isotopes average atomic mass is also known as Atomic weight (AW). Atomic weights are listed on the periodic table.

Naturally occurring Carbon is 98. 892 % 12C + 1. 108 % 13C Naturally occurring Carbon is 98.892 % 12C + 1.108 % 13C. Find the average atomic mass Step 1: Multiply the mass number of each isotope by its percentage Step 2: Add the numbers together Step 3: Divide by 100 Step 4: Check your answer. Mass should be between highest and lowest mass

An Isotope Example - Chlorine Let’s look at Chlorine, which has an atomic mass of 35.5 Atomic # = ___, which means that there are ___ protons. 17 17 There are two naturally occurring isotopes, Cl-35 and Cl-37 This means that Cl-35 has ___ neutrons while Cl-37 has ___ neutrons. 18 20 Cl-35 occurs naturally 75.8% of the time Cl-37 occurs naturally 24.2% of the time (35 x 75.8 + 37 x 24.2)/100 = 35.5

Try this one Calculate the average mass of the isotopes of Uranium with: 50.0% at 239 amu 29.4% at 235 amu 20.6% at 238 amu 237.62 amu

Classwork & Homework Activity Sheet on Parts of the Atom/Calculating Atomic Mass due Monday Lab Report due Wednesday September 10, 2014

Complete this chart Element/Ion Atomic Number Mass Number Protons Neutrons Electrons