Week 3 -Tuesday

FLAME TESTS, BOHR MODEL AND IONIC BONDS



Outline for the night

There is a lot of material but most of it should be review- for the most part 1.

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Review of energy levels in atoms / summary of the atom Flame test lab The modern periodic table A typical Atom – chlorine Bohr Diagrams Ions Lewis Structures Ionic Bonding

Review from last week

  Bohr suggested that electrons orbit the nucleus in fixed orbits of energy. Thus, the electrons are limited to certain energy levels and the energy of the electrons is quantized. Quantized – Possessing a specific value or amount (quantity)

Review from last week

 If energy is added and the electron moves up to a higher orbit it is said to be in an excited state. If the electron moves down to a lower orbit it must release the same amount of energy that was required to raise it. Therefore the lines spectra of hydrogen are the result of energy released when the hydrogen atoms electron falls to a lower state and releases energy.

Review from last week

 The lowest possible orbit for an electron is said to be its ground state. 

Flame Test Lab

   Activity 1.7 Page 23-24 Do not answer (h)

A Summary of the Atom

   All elements are composed of atoms, and one atom is the smallest unit of any element. The atoms themselves are made of different kinds of smaller particles, called

subatomic particles.

Three subatomic particles are protons, neutrons, and electrons, and they have different properties. One such property is relative mass.

A Summary of the Atom

   

Relative mass

compares the mass of an object to the mass of another object. An electron is the least massive subatomic particle of the three subatomic particles, so it is assigned a relative mass of 1. Compared to it, a proton has a relative mass of 1836, meaning that it is 1836 times heavier than an electron. Compared to an electron, a neutron is 1837 times heavier.

Name Proton Symbol Relative

p

Mass 1836 Electric Charge

1+

Location Nucleus Neutron Electron

n e

1836 1

0 1-

Nucleus In energy levels surround nucleus

Atoms are composed of three subatomic particles

1.

2.

3.

Protons Heavy positively charged found in the nucleus

Neutrons -

are neutral particles that have the same mass as protons and are located in the nucleus Electrons Negatively charged particles with almost no mass. They circle the nucleus at different energy levels.

Atoms are composed of three subatomic particles

 Atoms are electronically neutral so the number of electrons = the number of protons

THE MODERN PERIODIC TABLE

Dmitri Mendeleev’s Periodic Table (1834–1907)

The Modern Periodic Table

 The periodic table is a chart that places all of the elements in rows and columns. In the modern periodic table, elements are listed from left to right and top to bottom according to a property called

atomic number.

The Modern Periodic Table



Atomic Number (Z)

Atomic number -

the number of protons in an atom of an element.

  Each element has a set number of protons and every atom from that element will have that many protons. The pattern for increasing protons moves from left to right and then down to the next row just like reading a book.

Atomic Mass (A)

  Atomic mass is given in atomic mass units (amu).

 Atomic mass -

atoms. the average mass of an element’s

H has a mass of 1.01 amu. This means that iron atoms are about 55.85 times heavier than hydrogen atoms.

 Atomic masses are always expressed as decimal fractions. One reason that they do not have whole number values is that, except for fluorine, atoms of the same element have different numbers of neutrons.



Atomic Mass (A)

Example: A hydrogen atom has one proton and one electron but no neutron. A small percentage of hydrogen atoms have 1 p, 1e, and 1n. 1 p, 1e, and 2n.

  Atomic mass generally increases in order of atomic number. Exception: iodine (I) has a lower atomic mass than tellurium (Te).

Ion Charge

  An atom or group of atoms that has lost or gained electrons is called an ion. 

Ion charge the electric charge that an atom takes on when it loses or gains electrons.

Electrons have a negative charge, and so an atom that loses electrons becomes a positive ion. An atom that gains electrons becomes a negative ion.  Metal atoms can lose electrons in certain situations. (

positive ions

)  Non-metals can gain electrons in certain situations. (negative ions )

Ion Charge

 Elements with atoms that can form similar ions are grouped together in the periodic table. Metals generally lose electrons and become positive ions.  Many non-metals can gain electrons and so become negative ions.

Ion Charge

 Some elements do not form ions. Helium, for example, does not normally form ions. For these elements, no ion charges are shown in their squares in the periodic table.

Calculations

Determining the number of protons

 Look at the atomic number given on the periodic table (

atomic number)

 Example: H = 1, He = 2, Li = 3

Calculations

Determining the number of neutrons

 Subtract the atomic number (# of protons) from the Atomic mass (# of protons and neutrons) 

Atomic mass – atomic number (A – Z = N)

 Example: Iron 

55.85 – 26 = 30 neutrons

Calculations

Determining the charge of Ions

 Subtract the number of electrons from the number of protons   

Protons (P) – Electrons(E) = Ion Charge

Example: Iron  26 – 24 = 2+ charge =

Fe 2+

 26 – 23 = 3+ charge =

Fe 3+

Example: Fluorine 

9 – 10 = 1 – Charge = F -

Practice Questions–

           1. Use the periodic table to find the atomic number of each of the following elements.

(a) C (b) O (c) Na (d) Si (e) S (f) Cl (g) Fe 2. How many protons are in an atom of each of the following elements?

(a) lithium (b) nitrogen (c) fluorine (d) aluminum (e) copper (f) gold 3. Name the element with the following number of protons.

(a) 1 (b) 2 (c) 10 (d) 19(e) 20 (f) 31 (g) 47 4. Name the element with the following atomic mass.(a) 12.01 amu (b) 16.00 amu (c) 39.10 amu (d) 83.80 amu 5. What is the electric charge on an ion of each of the following elements?

(a) Li (b) Be (c) N (d) S (e) Al (f) I 6. Although the element hydrogen is a non-metal, it is located on the left side of the periodic table. Explain how placing hydrogen in this position relates to its ion charge.

7. Describe the patterns in atomic masses and ion charges in the periodic table.

Chlorine: A Typical Atom



Atoms of all elements have the same basic structure but different numbers of protons, neutrons, and electrons.

 The element chlorine is used as a disinfectant in swimming pools and to purify drinking water. A diagram of an atom of chlorine is shown below.  Notice that the number of protons and the number of electrons are equal. This is true of all atoms.

Example: Chlorine atom

   All chlorine atoms have 17 protons. Each proton has a charge of 1+, so the total positive charge in the nucleus is therefore 17+.

Different kinds of chlorine atoms can have different numbers of neutrons. It is the number of protons in an atom that determines what element the atom is, not the number of neutrons. The most common types of chlorine atoms have 18, 19, or 20 neutrons.

The Nucleus

  The nucleus is a tiny part of the atom that contains protons and neutrons gathered together into a ball. The nucleus contains only a small part of an atom’s total volume. Depending on the atom, the region outside the nucleus of an atom is 10 000 to 50 000 times the diameter of the nucleus.

The Nucleus

 The nucleus contains 99.99% of the mass of the atom because protons and neutrons have much greater mass than electrons

Electrons

Electrons exist in shells, or energy levels, surrounding the nucleus.

 The innermost shell can hold a maximum of two electrons. Each of the next two shells can hold up to eight. Electrons often exist in pairs in the shells.

 Electrons occupy more than 99.99% of an atom’s volume.

Electrons

 Electrons can move between energy levels.

 The outermost shell that has electrons in it is called the valence shell.  Electrons in this shell are called

valence electrons

.  Other shells containing electrons are called inner shells, and the electrons in them are called inner electrons.

Electrons

 

The properties of elements are strongly affected by their valence electrons.

When a shell becomes more than half-filled, the electrons begin to pair up, as shown in Bohr diagrams.  Even though the negatively charged electrons repel each other, pairing helps electrons to get closer to the positive protons in the nucleus.

– ELECTRON ARRANGEMENT

Patterns in the Arrangements of Electrons

 A very important pattern in the arrangement of electrons is that elements in the same group have the

same number of valence electrons

.

Patterns in the Arrangements of Electrons

 Group 1: Atoms of hydrogen, lithium, and sodium each have one valence electron. These elements share some chemical properties with them because of their similar valence electron arrangements. They all form ions with a

1+ charge

.

Patterns in the Arrangements of Electrons

 Group 18: Nobles gases have their valence shell completely filled with the

maximum number of electrons that they can hold.

 The noble gases share many properties because their atoms all have filled valence shells, this makes them very stable.

Patterns in the Arrangements of Electrons

 The number of valence electrons is not only related to the physical properties of a group of elements. The number of valence electrons is also related to the ways in which atoms of elements combine to form compound

Organizing the Periodic Table in Different Ways

Scientists continue to organize the elements in different ways.  Dr. Theodor Benfey, a U.S. chemist, suggested a spiral version of the periodic table.  In Dr. Benfey’s periodic table, the elements are shown in an unbroken series, starting with hydrogen and radiating outward.

Organizing the Periodic Table in Different Ways

 Another periodic table is known as the physicist’s periodic table. This periodic table is three dimensional and groups the elements according to the energy levels of their electrons. – It has been turned sideways to fit on the page

– BOHR DIAGRAMS

Bohr Diagrams

 To represent electron arrangements at various orbits we use Bohr diagrams. Each orbit has a set number of electrons.

Orbit # 1 2 3 4 # of Electrons

2 8 8 18

 Every row in the period contains a shell. The farther you move down the table the more shells you added to the diagram. H = 1 shell, Li = 2 shells, K = 3 shells.

 Moving left to right on the periodic table adds valence electrons to the shells of that row. Na has 1 valence e , Mg has 2 valence e , Al has 3 valence e , etc.

Drawing Bohr Diagrams

    Draw the Bohr diagram of Hydrogen N = A – Z N = N =

Drawing Bohr Diagrams

    Draw the Bohr diagram of Helium N = A – Z N = N =

Drawing Bohr Diagrams

    Draw the Bohr diagram of Lithium N = A – Z N = N =

Drawing Bohr Diagrams

    Draw the Bohr diagram of Beryllium N = A – Z N = N =

Drawing Bohr Diagrams

    Draw the Bohr diagram of Aluminum N = A – Z N = N =

Drawing Bohr Diagrams

    Draw the Bohr diagram of Argon N = A – Z N = N =

 Draw the Bohr diagrams for the first 20 elements. – Hand in

BOHR–RUTHERFORD MODEL: IONS

Bohr – Rutherford Model: Ions



Noble gases

do not form compounds because they have 8 electrons in their outer orbit (shell). This electron arrangement makes them very stable and so they do not react. N 10 P 10

 When elements form compounds, changes occur in the arrangement of electrons in the outer orbit. – Electrons are gained or lost so that element can have a stable electron arrangement of the closest noble gas. (In other words it will completely fill their outer shell with electrons)

 In order for a compound to be stable, it must have a completely filled outer electron shell– aka

(stable octet)



Arrangement of outer shell electrons of metals and non metals

Metals

  Tend to have 1, 2, or 3 electrons in the outer orbits (shells) They lose electrons when they combine with other elements to form positive ions (cations) : note the

t

in the word think

+

 They lose electrons, thus they have the same electron arrangement as the Noble gas a

row above

them

Metal Ion

Example  Sodium: Na  Na + N 12 P 11 N 12 P 11

Non-Metals

  Non-metals – Tend to have 4, 5, 6, or 7 electrons in their outer orbits (shells). They gain electrons to form negative ions (

anions

)  They gain electrons, thus they have the same electron arrangement as the Noble gas in the

same row

.

  Example Fluorine : F  F N 10 P 9 N 10 P 9

Homework

 Draw the ions of the first 20 ions of the periodic table

LEWIS SYMBOLS

Lewis Symbols

 Since it is only the valence shell electrons that take part in chemical reactions, a Lewis symbol drawing is used to depict the electrons in the atoms valence shell.

 Lewis Symbol –

A diagram composed of a chemical symbol and dots, depicting the valence electrons of an atom or ion

.

 When drawing Lewis structures, electrons are drawn on the four sides of the symbol. Each side can hold two electrons. The electrons are filled in one at a time until all the spaces are filled with one electron and then electrons are paired up. It’s like seats on the train. You don’t sit right beside someone if there are double empty seats all around you.

 Example of progressive placement- Oxygen

Lewis symbol Ions

 Draw the Lewis symbol for the ion showing all of the electrons in the outer shell. If there are none, just draw the symbol.  Then draw square brackets around the symbol and in the upper right corner write the charge on the ion.

THE FORMATION OF IONIC COMPOUNDS

Ionic Compounds



Ionic compounds

are formed by combining

metals

with

non-metals

in fixed proportions. An ionic compound is formed when one or more valence electrons are transferred from a metal atom to a non-metal atom. This leaves the metal ion as a cation and the non-metal ion as a anion

Ionic Compounds

 The two oppositely charged ions are attracted to each other by a force called a ionic bond.

Ionic Bond

 The smallest amount of substance that has the composition given by its chemical formula is the

formula unit.

 Ex.  NaCl 1:1 ratio  MgCl 2 1:2 ratio.

Ionic properties

 Ionic compounds are solids at SATP. In their solid form they form solid ionic crystals. These are more commonly known as salts

  This structure provides ionic compounds with the general properties of high MP, BP and hard. It takes a lot of energy to break the bonds between the ions. Any movement results in shift in where the ions are lined up and thus the like charges repel each other and the compound breaks.  This explains why many ionic compounds are brittle.

Conductivity

 When ionic compounds are dissolved in water they dissociate into their ions. This explains why ionic compounds conduct electricity when they are dissolved in water.  NaCl (s)  Na + (aq) + Cl (aq)

Conductivity

  The ions are able to carry the current, thus making dissolved ionic compounds conductive. This is called

electrolyte. Electrolyte

- A compound that when dissolved in water, produces a solution that conducts electricity.  A substance that when dissolved in water does not conduct electricity is called a non-

electrolyte.

Formation of an Ionic Bond

+ -

Lewis structure

 On the board

Practice Time

 Combine the following elements

Draw Bohr and Lewis structures

Li + F Al + N Mg + O Li + N H + F Na + S

Only Draw Lewis Structures

Fr + Se Ca + P Ba + As H + Te Ra + S Be + As Al + O

Ionic Bonding Video Clip

The Crossover rule

 This rule allows you to figure out how many atoms you will need of each element for bonding to occur without the need to draw Bohr diagrams

Crossover rule

Step 1.  Write the symbols, with the metal first (the element with the positive charge)  Mg I Step 2.  Write the Ionic charge above each symbol to indicate the stable ion that each element forms.  2+ 1  Mg I

Crossover Rule

    Step 3. Draw a arrow from the metals charge to the non-metal and an arrow from the non metal charge to the metal. (Cross over the arrows) 2+ Mg 1 I

Crossover rule

 Step 4. Fill in the number of atoms from each element will have by following the arrows. If need be reduce to lowest terms (in other words, if they are the same number, you don’t write those numbers down because you could divide the whole molecule by that number which would = 1)  MgI 2 (if the number crossed is a 1, the 1 is not shown)

          

Example 2



Ca + O

Step 1 Ca O Step 2 2+ 2 Ca O Step 3 2+ 2 Ca O Step 4 Ca 2 O 2 = CaO The 2’s disappeared because we reduced to lowest terms.

Practice

Li + Se Sr + O Al + Si Mg + P Rb+ N K + F Na +I Fr + N Ca + Br Ba + As H + Se Cs + N Be + Cl Al + O

Home Work

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3.

4.

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Hand in the lab Periodic table Practice Questions Bohr diagrams + Ion diagrams (first 20 on periodic table) Lewis Structures Ionic bonding practice