IGCSE CHEMISTRY 0620 Cambridge LESSON 2 EXPLANATION







  • Describe the build-up of electrons in shells and understand the significance of the noble gas electronic structures and of the outer shell electrons
 Electrons are arranged in shells around the nucleus in shells
Each shell has a different ability of holding the electrons 
The first shell can carry a maximum of 2 electrons 
when the atom has enough electrons to fill up this shell 
it starts to fill the second shell , which can carry a maximum of 8 , the third one carries 8 also . 

 The Noble gases are a group of elements which have a very special character , their outermost shell 
( the farthest one from the nucleus) is completely full with electrons , it can not carry anymore , it appears that these electronic configuration or arrangement gives the maximum stability for these atoms 

 That is why all of the other atoms want to achieve such arrangement by having full outer shell , so that they can attain the same stability 
For example The Hydrogen atom has 1 electron in its outermost shell , but the closest Noble gas element is Helium , which contains 2 electrons in its outermost shell (Full Shell) Helium is very stable element and it hardly reacts , But Hydrogen is very reactive , because it is always trying to fill its shell . 





  • 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.



Electron Shells

Electronic structure

  • We can represent the structure of the atom in two ways: using diagrams called electron shell diagrams or by writing out a special notation called the electronic configuration

Electron shell diagrams

  • Electrons orbit the nucleus in shells (or energy levels) and each shell has a different amount of energy associated with it
  • The further away from the nucleus then the more energy a shell has.
  • Electrons occupy the shell closest to the nucleus which can hold only 2 electrons and which go in separately
  • When a shell becomes full electrons then fill the next shell
  • The second shell can hold 8 electrons and the third shell can hold eighteen electrons and the electrons organise themselves in pairs in these shells
  • The outermost shell of an atom is called the valence shell and an atom is much more stable if it can manage to completely fill this shell with electrons

Electronic configuration

  • 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
  •  

  •  

 

Period: The red numbers at the bottom show the number of notations which is 3, showing that a chlorine atom has 3 shells of electrons

Group: The green box highlights the last notation which is 7, showing that a chlorine atom has 7 outer electrons

 


 

The noble gases

  • The atoms of the Group 8/0 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

 


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  • Describe the differences between elements, mixtures and compounds, and between metals and nonmetals
Substances can be very simple or very complicated according to how many kinds of atoms they are made of .
The simplest substances are made of one kind of atom only 
They are called elements 


Molecules are formed when two or more atoms are joined together ,they form what is known as a bond between them, for example two atoms of oxygen join together (bond) to form a molecule of oxygen

And now meet your first molecule


Atoms do not have to bond to the same kind of atoms only . in A molecule of water two atoms of hydrogen element join one atom of oxygen element to make a molecule of water , because this molecule is made of more than one kind of atom (element) we can not call it an element now , we call it a compound

Compounds are made when atoms make a chemical reaction together ,they form a totally new substance  that have different properties from the atoms they are made from

Example is sodium chloride table salt that we eat , can you imagine what are the properties of the elements this compound is made of 

Chlorine is a toxic gas , and sodium is a highly reactive metal 
can you eat a metal , can you put a highly toxic gas on you salad , yes only when they combine chemically to form a new compound that have properties totally different from the elements it was made from.



We can draw a molecule by drawing the atoms it is made of joined together


The colors shown here are only for showing the atoms better , you do not have to color your circles .


Sometimes the atoms mix together but without making a chemical reaction ( without making a new compound ) they will create something called a Mixture 

For Example in mud made of different compounds and elements (air , water , stones , rocks ) they are joined together without forming a new chemical compound or bonding together with chemical reactions , there is no chemical substance known as mud , only a mixture of physically joined substances , when you carry mud ,each component does not separate , sand does not fall down , or water out of the mixture . 
  • As a summary All substances can be classified into one of these three types.
Element
  • 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.
Compound
  • 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).
Mixture
  • 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.

Describing Alloys

  • Alloys are mixtures of metals, where the metals are mixed together but are not chemically combined
  • They can be made from metals mixed with nonmetals such as carbon
  • Alloys often have properties that can be very different to the metals they contain, for example they can have more strength, hardness or resistance to corrosion or extreme temperatures
  • Alloys contain atoms of different sizes, which distorts the regular arrangements of atoms
  • This makes it more difficult for the layers to slide over each other, so they are usually much harder than the pure metal
  • Brass is a common example of an alloy which contains 70% copper and 30% zinc

 Now the questions comes how the atoms fill up their shells ?
  • Describe the formation of ions by electron loss or gain
Ions
  • Unlike the atom an ion is an electrically charged atom or a group of atoms formed by the loss or gain of 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 electrostatic attraction 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:    NaCl

    Ionic 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

  • 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 repeating fashion.
    • The lattices formed by ionic compounds consist of a regular arrangement of 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.



    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

    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

    • 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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