Section 2 3 The Periodic Table and Atomic

Section 2. 3 – The Periodic Table and Atomic Theory Text Pages 60 -69

Periodic Table and Atomic Theory �The periodic table has a series of patterns! For example metals appear on the left and non-metals appear on the right.

Periodic Table and Atomic Theory �These patterns are no accident, and they result from changes in the structure of atoms as you move across the periodic table.

�Atomic Structure (Recap)

Recall �The atomic # for an element tells you the number of protons (p+) and electrons (e-) in an atom of that element

Recall � Electrons are found outside the nucleus of an atom. � Neils Bohr said that electrons exist in energy levels. *Electrons do not actually move in circles or ellipses. We have drawn it this way because it is the easiest way to represent it.

�In this section we will learn about the way that electrons are organized in energy levels and how this influences the reactivity of different elements.

STRUCTURE OF AN ATOM: ENERGY LEVELS

Energy Level: a region surrounding the nucleus of an atom that may be occupied by one or more electrons. They are also called electron shells.

Important Points about Energy Levels �The period # of an element tells you how many energy levels its atoms have. Ex: Carbon has ______ energy levels with electrons in them. Ex: Sulfur has ______ energy levels with electrons in them.

Important Points about Energy Levels �Energy levels can be numbered from 1 to 7. �Electrons in higher energy levels are further from the nucleus. �Electrons in higher energy levels have more energy.

The highest energy level in an atom that has electrons in it is known as the valence energy level. Any electrons in that level are known as valence electrons. Closer, Less Energy Further, More energy

�Not all energy levels can hold the same number of electrons. The rules below outline how many electrons each level can hold below. Energy Level Max. Number of Electrons 1 st Level 2 2 nd Level 8 3 rd Level 8 4 th Level 18

�Electrons enter these energy levels starting closest to the nucleus. �An energy level must be full before electrons start to occupy a higher level! �Assembly analogy

Bohr-Rutherford Diagrams �Bohr-Rutherford Diagram: shows the arrangement of the subatomic particles in an atom. It shows the number of protons and neutrons in the nucleus. It also shows the number of electrons in each energy level.

Drawing Bohr-Rutherford Diagrams �Step 1: Draw a circle to represent the nucleus. Write the element symbol inside at the top of the circle.

Drawing Bohr-Rutherford Diagrams �Step 2: Find atomic number gives # of electrons (e-) and # of protons (p+)

Drawing Bohr-Rutherford Diagrams �Step 3: Calculate # of neutrons (no) using Mass # - Atomic #

Drawing Bohr-Rutherford Diagrams �Step 4: Label inside the nucleus with # of p+ AND # of no (write these below the element symbol!)

Drawing Bohr-Rutherford Diagrams �Step 5: Fill the energy levels using the 2, 8, 8, 18 energy level pattern for electrons. Each energy level looks like a dashed horizontal line with # of ewritten inside it: --5 e--

Worksheet #11 �Complete elements. diagrams for the first 18

Remember:

Valence Electrons and Reactivity � Do you notice any relationship between the maximum # of electrons in an energy level and the # of elements in each period?

What do you notice about their electron structures? �Valence energy levels? �# of valence electrons?

Elements in the same family (group) have the same number of valence electrons. This explains why they have similar properties.

�The number of energy levels in the atoms of an element increases as you move down a group.

�Elements in the same period have electrons in the same valence energy level.

Valence electrons are very important because the number of valence electrons determines how reactive an element will be.

�If an element has a full valence energy level, it will not be reactive. We say that it is stable. This helps explain why the Noble Gases are stable. �Example: Neon

�Elements that do not have full valence energy levels will react with other elements in order to become stable! These reactions involve gaining, losing, or sharing electrons! �Examples: Lithium and Fluorine