Atomic Structure, Bonding, and Properties

Atomic Structure and the Periodic Table

The Periodic Table, as developed by Henry Moseley, organizes elements based on their atomic structure. An atom consists of three subatomic particles: protons, neutrons, and electrons. Protons carry a positive charge ($+1$), electrons carry a negative charge ($-1$), and neutrons have no charge ($0$). An atom has no overall charge because the number of protons is identical to the number of electrons.

Every element is defined by its atomic number, which tells you exactly how many protons it contains. The atomic number increases by one for every sequential element in the Periodic Table. Additionally, every element has a mass number, which represents the total sum of protons and neutrons in the nucleus. For example, Lithium ($Li$) has an atomic number of $3$ and a mass number of $7$. To determine the number of neutrons, you subtract the atomic number from the mass number: $7 - 3 = 4\, \text{neutrons}$.

The atoms of the elements increase in mass as you progress from left to right and downwards in the Periodic Table. For example, an atom of hydrogen has significantly less mass than an atom of sodium. The Periodic Table categorizes elements into metals and non-metals. Metals like lithium ($Li$), beryllium ($Be$), sodium ($Na$), magnesium ($Mg$), potassium ($K$), and calcium ($Ca$) are generally found on the left and center, while non-metals like boron ($B$), carbon ($C$), nitrogen ($N$), oxygen ($O$), fluorine ($F$), neon ($Ne$), phosphorus ($P$), sulfur ($S$), chlorine ($Cl$), and argon ($Ar$) are on the right.

Properties of Materials and Density

The mass of an atom is directly correlated to its mass number; for instance, because Carbon has an atomic mass of $12$ and Hydrogen has an atomic mass of $1$, one atom of carbon has $12$ times more mass than one atom of hydrogen. This physical property leads to the concept of density, which is defined as the measure of mass for a fixed volume of a substance. The standard unit for density is $g/cm^3$.

In a comparison of fixed volumes ($1\, cm^3$), Sodium ($Na$) has a mass of $0.97\, g$, resulting in a density of $0.97\, g/cm^3$. In contrast, Iron ($Fe$) has a mass of $7.87\, g$, resulting in a density of $7.87\, g/cm^3$. Therefore, iron is significantly more dense than sodium.

Electronic Structure and the Bohr Model

The modern understanding of the atom was built upon the models of J. J. Thompson and Ernest Rutherford. In 1913, the Danish scientist Niels Bohr advanced Rutherford's model by proposing that electrons move in different electron shells, also known as energy levels, around the nucleus. For this revolutionary work, Bohr was awarded the Nobel Prize.

The electronic structure refers to the specific arrangement of electrons in these shells. The first electron shell is closest to the nucleus and has a capacity for only $2$ electrons. The second and third electron shells have a larger capacity, holding up to $8$ electrons each. Electrons are maintained in their positions by electrostatic forces of attraction between the negatively charged electrons and the positively charged nucleus.

Electronic arrangements are often written as numbers separated by commas. For example, Boron ($B$) has an electronic structure of $2,3$, indicating two electrons in the first shell and three in the second. The first shell must always be completely filled before electrons occupy the second shell. A modern refinement of this model notes that electrons are not in fixed orbits but exist in areas of probability based on energy levels.

Group 1: The Alkali Metals

Group 1 of the Periodic Table consists of the alkali metals, including Lithium ($Li$), Sodium ($Na$), and Potassium ($K$). These elements share specific structural and chemical properties. As you move down Group 1, both the atomic number and mass number increase, indicating that the size of the atom is increasing.

There are distinct trends in physical properties for Group 1:

• Lithium ($Li$): Atomic Number $3$, Mass Number $7$, Electronic Structure $2,1$, Melting Point $180\, ^{\circ}C$, Boiling Point $1360\, ^{\circ}C$.
• Sodium ($Na$): Atomic Number $11$, Mass Number $23$, Electronic Structure $2,8,1$, Melting Point $98\, ^{\circ}C$, Boiling Point $900\, ^{\circ}C$.
• Potassium ($K$): Atomic Number $19$, Mass Number $39$, Electronic Structure $2,8,8,1$, Melting Point $63\, ^{\circ}C$, Boiling Point $777\, ^{\circ}C$.

The trend shows that melting and boiling points decrease as you go down the group. This is because metallic bonding—the force that bonds a metal atom to the free electrons present in the metallic structure—becomes weaker as atoms get larger and the nuclear pull on the outermost electron decreases. Based on these trends, one can predict that Rubidium ($Rb$), the next element in the group, would have a melting point lower than $63\, ^{\circ}C$.

Group 7: The Halogens

Group 7 (or Group 17 in some numbering systems) is known as the Halogens. This group includes Fluorine ($F$), Chlorine ($Cl$), and Bromine ($Br$). These elements are highly reactive non-metals and represent all three states of matter at room temperature: Fluorine and Chlorine are gases, while Bromine is a liquid.

Key data for Group 7 elements:

• Fluorine ($F$): Atomic Number $9$, Mass Number $19$, Electronic Structure $2,7$, Color: Pale yellow, Melting Point $-220\, ^{\circ}C$, Boiling Point $-188\, ^{\circ}C$.
• Chlorine ($Cl$): Atomic Number $17$, Mass Number $35$, Electronic Structure $2,8,7$, Color: Yellowish-green, Melting Point $-101\, ^{\circ}C$, Boiling Point $-34\, ^{\circ}C$.
• Bromine ($Br$): Atomic Number $35$, Mass Number $80$, Color: Brown, Melting Point $-7\, ^{\circ}C$, Boiling Point $59\, ^{\circ}C$.

In Group 7, the reactivity decreases down the group; Fluorine is the most reactive, and Bromine is less reactive than chlorine. Conversely, the melting and boiling points increase as you go down the group, and the color of the elements becomes darker. All halogens have seven electrons in their outermost shell.

Group 8: The Noble Gases

Group 8 (or Group 0/18) consists of the Noble Gases, including Helium ($He$), Neon ($Ne$), and Argon ($Ar$). These elements are gases at room temperature and are characterized by being inert (unreactive) because their outermost electron shells are completely full.

Data for Group 8 elements:

• Helium ($He$): Atomic Number $2$, Mass Number $4$, Electronic Structure $2$, Melting Point $-270\, ^{\circ}C$, Boiling Point $-269\, ^{\circ}C$.
• Neon ($Ne$): Atomic Number $10$, Mass Number $20$, Electronic Structure $2,8$, Melting Point $-249\, ^{\circ}C$, Boiling Point $-246\, ^{\circ}C$.
• Argon ($Ar$): Atomic Number $18$, Mass Number $40$, Electronic Structure $2,8,8$, Melting Point $-189\, ^{\circ}C$, Boiling Point $-186\, ^{\circ}C$.

Because they have full outer shells, noble gases are extremely stable and do not naturally form compounds. The size of the atoms increases as you move down the group.

Chemical Bonding and Ion Formation

Atoms react to form compounds to achieve stability, which occurs when their outermost electron shell (highest energy level) is completely full. Chemical bonds are formed either by losing or gaining electrons (ionic bonding) or by sharing electrons (covalent bonding).

An ion is a charged particle formed when an atom loses or gains electrons. Metals typically lose electrons to form positive ions ($cations$). For example, a Sodium atom ($2,8,1$) loses its one outer electron to become a Sodium ion ($Na^+$) with a structure of $[2,8]^+$. This ion is more stable because its new outermost shell is full.

Non-metals typically gain electrons to form negative ions ($anions$). For example, a Chlorine atom ($2,8,7$) gains one electron to become a Chloride ion ($Cl^-$) with a structure of $[2,8,8]^-$.

Ionic bonding is the electrostatic attraction between these oppositely charged ions. In Sodium Chloride ($NaCl$), the sodium atom gives one electron to the chlorine atom. Similarly, Magnesium ($2,8,2$) can react with Oxygen ($2,6$) to form Magnesium Oxide ($MgO$). The magnesium atom loses two electrons to become $Mg^{2+}$, and the oxygen atom gains those two electrons to become $O^{2-}$.

Reactivity Trends Explained

The reactivity of Group 1 metals increases as you go down the group. This is because the single electron in the outermost shell of Potassium ($K$) is further from the positive nucleus than the outer electron in Lithium ($Li$). The electrostatic force holding the electron is weaker at a greater distance, making it easier for Potassium to lose its electron and react.

In contrast, the reactivity of Group 7 halogens decreases as you go down the group. Fluorine is more reactive than Chlorine because its outermost shell is closer to the nucleus. This creates a stronger electrostatic pull, allowing the Fluorine atom to attract and gain an additional electron more easily than a Chlorine atom.

Covalent Bonding

Covalent bonds form when non-metal atoms share pairs of electrons to achieve full outer shells. This results in the formation of molecules. A common way to represent these is through "dot and cross diagrams."

Examples of covalent molecules include:

• Hydrogen molecule ($H_2$): Two hydrogen atoms share one pair of electrons.
• Hydrogen Chloride ($HCl$): One hydrogen atom and one chlorine atom share one pair of electrons.
• Water ($H_2O$): Two hydrogen atoms each share a pair of electrons with one oxygen atom.
• Ammonia ($NH_3$): Three hydrogen atoms share three pairs of electrons with one nitrogen atom ($N_{atomic\, no.}=7$, structure $2,5$).
• Methane ($CH_4$): Four hydrogen atoms share electrons with one carbon atom.
• Other molecules include Oxygen ($O_2$), Nitrogen ($N_2$), Carbon Dioxide ($CO_2$), and Chlorine ($Cl_2$).

Simple and Giant Structures

Substances can form simple molecular structures or giant structures. Ionic compounds like Sodium Chloride form a giant lattice structure. In an $NaCl$ lattice, each $Na^+$ ion is surrounded by six $Cl^-$ ions, creating a regular repeating pattern that results in crystals with a regular shape. These structures are held together by strong electrostatic forces acting in all directions.

Most covalent substances form simple molecules (like $CO_2$ or $CH_4$). While the covalent bonds within the molecules are strong, the intermolecular forces between molecules are weak. Consequently, simple covalent molecules have low melting and boiling points.

Some covalent substances form giant covalent structures (macromolecules):

• Diamond: Each carbon atom forms four strong covalent bonds in a rigid, three-dimensional lattice. This makes diamond the hardest natural material, used in jewelry and industrial drill bits.
• Graphite: Each carbon atom bonds to three others, forming layers. The bonds within layers are strong, but the bonds between layers are weak, allowing them to slide. This makes graphite soft and useful as a lubricant and for pencil "lead."
• Silicon Dioxide ($SiO_2$): The main component of sand, which has a high melting point ($1610\, ^{\circ}C$) due to its giant covalent structure.

Physical and Electrical Properties

The type of bonding determines the physical properties of a substance:

• Melting and Boiling Points: Ionic substances (like $NaCl$, $MgCl_2$, $CaO$) have very high melting and boiling points due to strong electrostatic forces. Simple covalent molecules (like $CH_4$, $NH_3$, $Cl_2$) have low melting and boiling points due to weak intermolecular forces.
• Electrical Conductivity: Ionic compounds conduct electricity only when dissolved in water (aqueous) or melted (molten). This is because the ions, which carry the electrical charge, are free to move. In solid form, they do not conduct. Simple covalent substances do not conduct electricity because they do not have free-moving charged particles.

Questions & Discussion

1. How many electrons are there in an atom of carbon? An atom of carbon has $6$ electrons.

2. How many protons are there in an atom of beryllium? There are $4$ protons in an atom of beryllium.

3. How many neutrons are there in an atom of boron? Boron has a mass number of $11$ and an atomic number of $5$. Therefore, it has $11 - 5 = 6\, \text{neutrons}$.

4. What is the name of the element that has the electronic structure 2,8,3? Aluminium.

5. How is the model of the atom we use today different from the Rutherford model? The modern atomic model differs because electrons are not in fixed orbits; they move in energy levels around the nucleus, described as a probability density.

6. What is the trend in the boiling points of Group 1 elements? The boiling points decrease as you move down the group ($1360\, ^{\circ}C$ for $Li$ to $777\, ^{\circ}C$ for $K$).

7. Suggest why the next element, Rubidium, is not used in schools. Rubidium is extremely reactive (more so than potassium) and can react explosively with moisture in the air or water, making it too dangerous for standard school demonstrations.

8. Why do the items magnesium chloride and calcium oxide have high melting points? They are ionic compounds held together by strong electrostatic forces between ions, which require a large amount of energy to overcome.