Home CIE IGCSE Chemistry Lesson 2
1 Describe the build up of electrons in shells and understand the significance of noble gas electronic structures and of the outer shell electrons . (The ideas of the distribution of electrons in s and p orbitals and in d block elements are not required.)
1 Describe the build up of electrons in shells and understand the significance of noble gas electronic structures and of the outer shell electrons . (The ideas of the distribution of electrons in s and p orbitals and in d block elements are not required.)
Electrons are arranged in shells around the nucleus
- The electrons will fill up the shell closet to the nucleus first (shell 1 ) then shell 2 and shell 3 and so on
Example Electron arrangement in the atom of Argon
So the arrangement or the electronic structure of an atom of argon has
2 electrons in shell 1
8 In shell 2
8 in shell 3
All the elements that exists Naturally or synthetically are arranged in the periodic table
the Periodic table starts with the element with 1 proton (Hydrogen) and then next lies the element with 2 protons and so on
The elements in the periodic table are lined in horizontal and vertical lines
we call the Horizontal lines (PERIODS) and the Vertical lines (Groups) ,elements with similar chemical properties are arranged in the same group
- The arrangement of electrons in shells can also be explained using numbers.
- There is a clear relationship between the outer shell electrons and how the Periodic Table is designed.
- The number of notations in the electronic configuration will show the number of shells of electrons the atom has, showing the Period in which that element is in.
- The last notation shows the number of outer electrons the atom has, showing the Group that element is in.
- Elements in the same Group have the same number of outer shell electrons.
For Example in Atom of Chlorine
its electronic configuration is as follows
2 , 8 , 7
so it has 3 notations , which means it has 3 shells , and that it is in PERIOD number 3 in periodic table , the last number which is 7 shows it has 7 electrons in its last shell ( called the outermost shell ) and that the element lies in GROUP 7 in the periodic table
The noble gases (what is so special about them)
- The atoms of the Group 8 sometimes called 0 elements all have 8 electrons in their outer shells, with the exception of helium which has 2. But since helium has only 2 electrons in total and thus the first shell is full (which is the only shell), it is thus the outer shell so helium also has a full valency shell.
- All of the noble gases are unreactive as they have full outer shells and are thus very stable.
- All elements wish to fill their outer shells with electrons as this is a much more stable and desirable configuration.
Note: although the third shell can hold up to 18 electrons, the filling of the shells follows a more complicated pattern after potassium and calcium. For these two elements, the third shell holds 8 and the remaining electrons (for reasons of stability) occupy the fourth shell first before filling the third shell.
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Types of Substance and Properties
Describe the differences between elements, mixtures and compounds, and between metals and nonmetals
Elements, compounds and mixtures
- All substances can be classified into one of these three types.
- A substance made of atoms that all contain the same number of protons (one type of atom) and cannot be split into anything simpler.
- There is a limited number of elements and all elements are found on the Periodic Table.
- E.g. hydrogen, carbon, nitrogen.
- A pure substance made up of two or more elements chemically combined together.
- There is an unlimited number of compounds.
- Compounds cannot be separated into their elements by physical means.
- E.g. copper (II) sulphate (CuSO4), calcium carbonate (CaCO3), carbon dioxide (CO2).
- A combination of two or more substances (elements and/or compounds) that are not chemically combined.
- Mixtures can be separated by physical methods such as filtration or evaporation.
- E.g. sand and water, oil and water, sulphur powder and iron filings.
Metals and nonmetals
- The Periodic Table contains over 100 different elements.
- They can be divided into two broad types: metals and nonmetals.
- Most of the elements are metals and a small number of elements display properties of both types. These elements are called metalloids or semimetals.
Properties of metals
- Conduct heat and electricity.
- Are malleable and ductile (can be hammered and pulled into different shapes).
- Tend to be lustrous (shiny).
- Have high density and usually have high melting points.
- Form positive ions through electron loss.
- Form basic oxides.
Properties of nonmetals
- Do not conduct heat and electricity.
- Are brittle and delicate when solid and easily break up.
- Tend to be dull and nonreflective.
- Have low density and low melting points (many are gases at room temperature).
- Form negative ions through electron gain (except for hydrogen).
- Form acidic oxides.
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The Formation of Ions
Describe the formation of ions by electron loss or gain
ons
- An ion is an electrically charged atom or a group of atoms formed by the loss or gainof electrons.
- This loss or gain of electrons takes place to gain a full outer shell of electrons.
- The electronic structure of an ion will be the same as that of a noble gas – such as helium, neon and argon.
Ionisation of metals and non-metals
- Metals: all metals lose electrons to another atom and become positively charged ions.
- Non-metals: all non-metals gain electrons from another atom to become negatively charged ions.
Electrostatic attraction
- The positive and negative charges are held together by the strong electrostaticattraction between opposite charges.
- This is what holds ionic compounds together.
Core:
Describe the formation of ionic bonds between elements from Groups I and VII
Example: Sodium Chloride, NaCl
Explanation- Sodium is a group 1 metal so will lose one outer electron to another atom to gain a full outer shell of electrons.
- A positive ion with the charge +1 is formed.
- Chlorine is a group 7 non-metal so will need to gain an electron to have a full outer shell of electrons.
- One electron will be transferred from the outer shell of the Sodium atom to the outer shell of the Chlorine atom.
- Chlorine atom will gain an electron to form a negative ion with charge -1.
Formula of Ionic Compound: NaClIonic Bonds between Metallic and Non-Metallic Elements
Describe the formation of ionic bonds between metallic and non-metallic elements
Example: Magnesium Oxide, MgO
Diagram Showing the Dot-and-Cross Diagram of Magnesium Oxide
Explanation
- Magnesium is a group 2 metal so will lose two outer electrons to another atom to have a full outer shell of electrons.
- A positive ion with the charge +2 is formed.
- Oxygen is a group 6 non-metal so will need to gain two electrons to have a full outer shell of electrons.
- Two electrons will be transferred from the outer shell of the Magnesium atom to the outer shell of the Oxygen atom.
- Oxygen atom will gain two electrons to form a negative ion with charge -2.
Formula of ionic compound: MgO
Supplement:
Describe the lattice structure of ionic compounds as a regular arrangement of alternating positive and negative ions
- Lattice structure refers to the arrangement of the atoms of a substance in 3D space.
- In these structures the atoms are arranged in an ordered and repeatingfashion.
- The lattices formed by ionic compounds consist of a regular arrangementof alternating positive and negative ions.
Core:
Describe the formation of single covalent bonds in H2, Cl2, H2O, CH4, NH3 and HCl as the sharing of pairs of electrons leading to the noble gas configuration
Covalent compounds- Covalent compounds are formed when electrons are shared between atoms.
- Only non-metal elements participate in covalent bonding.
- As in ionic bonding, each atom gains a full outer shell of electrons.
- The bonded unit of atoms is what is called a molecule.
Core:
Describe the differences in volatility, solubility and electrical conductivity between ionic and covalent compounds
- Ionic compounds:
- Have high melting and boiling points so ionic compounds are usually solid at room temperature.
- Not volatile so they don’t evaporate easily.
- Usually water soluble as both ionic compounds and water are polar
- Conduct electricity in molten state or in solution as they have ions that can move and carry charge.
- Covalent compounds:
- Have low melting and boiling points so covalent compounds are usually liquidsor gases at room temperature.
- Usually volatile which is why many covalent organic compounds have distinct aromas.
- Usually not water soluble as covalent compounds tend to be nonpolar but can dissolve in organic solvents.
- Cannot conduct electricity as all electrons are involved in bonding so there are no free electrons or ions to carry the charge.
Electron Arrangement in Complex Covalent Molecules
Supplement:
Describe the electron arrangement in more complex covalent molecules such as N2, C2H4, CH3OH and CO2
Melting and Boiling Points of Ionic and Covalent Compounds
Supplement:
Explain the differences in the melting point and boiling point of ionic and covalent compounds in terms of attractive forces
- Ionic compounds have high melting and boiling points.
- This is because the ions in the lattice structure are attracted to each other by strong electrostatic forces which hold them firmly in place.
- Large amounts of energy are needed to overcome these forces so the m.p. and b.p. are high.
- Covalent substances have very strong covalent bonds between the atoms, but much weaker intermolecular forces holding the molecules together.
- When one of these substances melts or boils, it is these weak intermolecular forces that break, not the strong covalent bonds.
- So less energy is needed to break the molecules apart so they have lowerm.p. and b.p than ionic compounds.
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