Questions:

• Atomic radius versus atomic number

A number of physical and chemical properties are related to the sizes of the atoms, but atomic size is somewhat difficult to define. There is no precise outer boundary of an atom. The radius is one half the distance between the centers of two adjacent atoms. The radius of an atom depends on the environment in which it is found. For bonded atoms, we customarily speak of a covalent radius, ionic radius, and, in the case of metals, a metallic radius. For atoms that are not bonded together, the radius is known as the van der Waals radius. For comparison, all radii in the above table are covalent.

1. 1. Which is the largest of the first 54 elements?

Rubidium (at 0.216 nm)

1. 1. Describe how the atomic radius varies within a period and within a family.

In a general description you can determine that atomic radii increases as you go from the top of your family to your bottom and as you go from right to left along a family generally, because as you might think as your number of protons increases so would your radii size, but no. As your number of protons goes up so does the attraction between the electrons pulling them tighter together so it actually gets more compressed.

1. 1. Use your graph to predict the atomic radius of the following elements:
      1. cesium  0.265 (b)   tungsten 0.145 (c) thallium  0.169 (d)   radon 0.140
   2. Which group of the main group elements contains the largest elements?

The group that contains the largest elements (when referring to atomic radii size) is group 1, this is because their valence shells contain one electron and that one electron pull is lessened by the full shell before (it’s hard for the electron to overcome, sort of “blocked out”), and also because it would take less energy to be unstable.

• Ionization energy versus atomic number

1. 1. How would you explain ionization energy to your partner?

I would explain ionization energy to my partner like this: Ionization energy is the energy needed to take away  valence electron(s) from an atom. It is also highly dependent on the atomic radii, because the farther away your electrons are from your nucleus the easier it is for them to get “stolen” they can be considered the looser electrons.

1. 1. 1. How does the ionization energy vary within a period and within a family?

Moving across the period, the ionization energy increases because the attraction of electrons to the nucleus also increases. In a family, when moving downwards the ionization energy decreases.

1. 1. 1. Which element on your graph has the strongest hold of its valence electrons? Helium.
   2. (a) Write the electron configuration for chlorine. Cl: 1s2 2s2 2p6 3s2 3p5

(b) Which electron is lost when 1251 kJ/mol of energy are applied to a sample of chlorine atoms? 3p5 because 1251 kJ/mol is the amount of energy it takes to remove an electron from a neutral chlorine atom.

2. Compare the ionization energies of metals to nonmetals.

Nonmetals and metals ionization can be compared to their atomic radii, because non metals valence electrons have stronger connections to their nucleus it takes more energy to pull them away, so instead they accept electrons to become negative ions. However metals have generally low ionization energies making it less work to take electrons from their valence shell, forming positive ions.

• Melting point versus atomic number

1. 1. Describe the trend of melting points within a period

The strength of a metallic bond causes the melting point to increase among the periods. The atoms will minimize and get smaller as you move left to right on the periodic table.

1. 1. Which group of elements tends to have the highest melting points

The metals have the highest melting points because their bonds are so strong it requires more energy to change the state of the element.

1. 1. Tungsten is used in incandescent light bulbs because it has an extremely high melting point. Which element on your chart could be a reasonable replacement for tungsten? Why?

Carbon could potentially replace tungsten, because it has the highest melting point within the entire periodic table.

• Density versus atomic number

1. 1. Describe how density varies within a period.

Often the density increases when you go top to bottom (of the family) and left to right (of the period), although solid elements will be higher than gaseous elements.

1. 1. Compare the densities of the elements in the second period with the elements in the third period.

From the second period to the third period, the density increases as you move further right.

1. 1. Assume that the transition metals given in the table are representative of the other members of this group. How do the densities of the transition metals compare with those of the elements in the main  group?Out of every group on the periodic table, the transition metals have the highest density. The elements in the main group are not as dense as the transition metals.
      1. 1. Explain why aluminum and magnesium are more suitable than iron for use in some airplane parts.

      Magnesium and Aluminum are more suitable than iron in some airplane parts because they are both lightweight, but yet durable. Iron has a higher density than Magnesium and Aluminum (almost 4x as dense).

   2. Electronegativity versus atomic number

   3. 1. 1. Describe how electronegativity varies within a period.

      As you move left to right in each period, the electronegativity increases for each element.

      1. 1. Describe how electronegativity varies within a family.

      As you move from top to bottom in a family, the electronegativity will decrease. This occurs because the valence electrons are further away from the nucleus.

      
      

      Graph A:

      Graph B:

      Graph C:

      Graph D:

      Graph E: