How do you find the valence of Mg?

The valence of an element refers to the number of bonds an atom of that element can form. To find the valence of magnesium (Mg), we need to look at its position on the periodic table and its electron configuration.

Quick Answers

– The valence of magnesium is 2.
– Magnesium has 2 valence electrons in its outermost shell.
– Magnesium commonly forms ionic compounds by losing 2 electrons to form a Mg2+ ion.

On the periodic table, magnesium is located in group 2, also known as the alkaline earth metals. Elements in this group have a valence of 2. This means that magnesium atoms will tend to lose 2 electrons in chemical reactions to form ions with a 2+ charge.

We can confirm this by looking at the electron configuration of magnesium. Magnesium has an atomic number of 12, meaning it has 12 protons in its nucleus. A neutral magnesium atom will also have 12 electrons.

The electron configuration of magnesium is: 1s2 2s2 2p6 3s2

The highest energy level, or outermost shell, is the third level, which contains 2 electrons. These 2 electrons in the outer shell are called valence electrons. Magnesium readily loses these 2 valence electrons to achieve a more stable octet configuration.

Determining Valence from Electron Configuration

In general, the valence of an element can be determined by looking at the number of electrons in its outermost energy level. Elements in groups 1, 2, and 13-18 are more likely to lose, gain, or share electrons to achieve a full octet in their valence shell.

For example:

  • Group 1 elements have 1 valence electron and a valence of 1
  • Group 2 elements have 2 valence electrons and a valence of 2
  • Group 13 elements have 3 valence electrons and a valence of 3

Transition metals and other elements may have more variable valences depending on their bonding situation. However, the main group elements tend to adhere to the pattern of their group number being equal to their typical valence.

Magnesium Forming Ionic Bonds

Magnesium readily forms ionic bonds by losing its 2 valence electrons to become a cation with a +2 charge. Some common magnesium compounds include:

  • Magnesium oxide (MgO) – Mg2+ and O2- ions
  • Magnesium chloride (MgCl2) – Mg2+ and Cl- ions
  • Magnesium sulfate (MgSO4) – Mg2+ and SO42- ions

In these ionic compounds, the magnesium cation has lost 2 electrons and has a +2 charge. We can see that magnesium obeyed the common valence rule, losing electrons equal to its group number to become stable.

Magnesium in Covalent Compounds

While less common, magnesium can also form covalent bonds, sharing its two valence electrons to obtain a full octet. For example, magnesium can react with hydrogen gas to form magnesium hydride (MgH2).

In MgH2, magnesium forms two covalent bonds, sharing its valence electrons in bonded pairs with the hydrogen atoms.

Even in covalent compounds, we see that magnesium utilizes its two available valence electrons to satisfy its bonding requirements.

Determining Variable Valences

Transition metals can exhibit variable valence states based on their electron configuration. However, the most common valence can still be deduced from the electron configuration.

For example, copper has an electron configuration of [Ar] 3d10 4s1. The valence shell is the 4s subshell with 1 electron. Therefore, copper will commonly exhibit a valence of 1, forming ions such as Cu+ and Cu2+ by losing electrons from its partially filled d subshell.

Other transition metal valences can be determined similarly by identifying the number of s and d electrons present in the outermost shell.


In summary, the valence of magnesium and other elements can be determined by:

  • Identifying the group on the periodic table
  • Examining the electron configuration
  • Noting the number of electrons in the outermost shell

Using these steps, we find magnesium has a valence of 2, owing to its position on the table, its stable electron configuration of Ne 3s2, and its two valence electrons that it readily gives up to form cations like Mg2+. Determining valence allows us to predict the typical bonding behavior and reactivity of the elements.

Element Electron Configuration Valence Electrons Valence
Magnesium 1s2 2s2 2p6 3s2 2 2
Sodium [Ne] 3s1 1 1
Calcium [Ar] 4s2 2 2

This table summarizes the electron configuration, valence electrons, and common valence for magnesium and other alkaline earth metals.

Finding Valence Using Periodic Trends

In addition to the steps outlined above, periodic trends can also help determine valence:

  • Metals tend to lose valence electrons to form cations
  • Nonmetals tend to gain electrons to form anions
  • The charge of transition metal ions increases across a row as electrons fill d subshells

Using position on the periodic table as a guide, metals on the left side tend to have lower valences of 1-3, while nonmetals on the right side tend to have valences of -1 to -3. Transition metal valences get higher across the rows.

Variable Valence States

Some elements exhibit multiple valence states under different conditions. However, there is often one predominant state:

  • Iron – Common states are 2+ and 3+
  • Tin – Common states are 2+ and 4+
  • Lead – Common states are 2+ and 4+

Variable valences are observed when elements undergo oxidation-reduction reactions and change valence state. But examining electron configurations and periodic trends can reveal the most likely valence state.

Using Valence to Determine Molecular Formulas

Knowing the valence of elements allows determination of unknown molecular formulas. For example, an ionic compound is made of Na+ and O2-. Sodium (Na) has a valence of 1+, while oxygen (O) has a typical valence of 2-. Using these valences, we can determine the formula is Na2O.

For covalent molecules, valence can also predict bonding patterns. Phosphorus (P) has a valence of 5, nitrogen (N) has a valence of 3. Therefore, phosphorus pentachloride (PCl5) and diphosphorus pentoxide (P2O5) satisfy the valence requirements.

Determining valence is key for balancing chemical equations, naming compounds, comparing reactivity, understanding periodic trends, and numerous other applications.

Valence vs Oxidation States

Although sometimes used interchangeably, valence and oxidation states are slightly different:

  • Valence – number of bonds an atom can form (doesn’t have to equal charge)
  • Oxidation state – charge an atom would have if all bonds are fully ionic

For main group elements, the valence often equals the oxidation state. However, for transition metals with partly covalent bonds, the oxidation state may differ from the valence.

For example, in H2SO4 the sulfur has a valence of 6 from forming six covalent bonds. But the formal oxidation state of S is +6 assumingionic character. The valence reflects its bonding behavior, while oxidation state assumes complete ionicity.

Other Ways to Determine Oxidation States

Rules for assigning oxidation states include:

  • The oxidation state of an uncombined element is zero (e.g. Na0, O0)
  • The sum of oxidation states in a compound is zero
  • The sum of oxidation states in a polyatomic ion is equal to the charge
  • Fluorine is always -1 unless bonded to oxygen or another fluorine
  • Group 1 metals are always +1, group 2 metals are always +2, etc.

Using these rules can help determine oxidation state when it may differ from the element’s valence.

Importance of Valence

Understanding elemental valence is critically important across many fields including:

  • Chemistry – predicting reactivity, formulas, modeling bonding
  • Materials Science – relating properties to valence electron configurations
  • Engineering – designing compounds and alloys with desired properties
  • Biology – appreciating bonding in biomolecules and cofactors
  • Geology – modeling formation of minerals and earth materials

Valence electron patterns give rise to periodic trends in reactivity. Metals tend to lose valence electrons, while nonmetals tend to gain electrons. Valence dictates preferred bonding patterns and molecular geometries.

Creative utilization of elemental valence states allows the design of advanced materials like semiconductors and superconductors. Understanding valence is foundational knowledge for many scientific disciplines.

Real-World Applications of Valence

Some examples of applying valence understanding include:

  • Batteries – cathode and anode materials utilize controlled valence states to drive electrons through the circuit
  • Catalysts – tailored valence states activates sites for chemical reactions, like nitrogen fixation
  • Photovoltaics – semiconductors with known valence states convert solar energy to electricity
  • Vitamins/Cofactors – utilize transition metals with variable valence to transfer electrons in biological processes

Rational design relies on mastering the valence configurations and tendencies of the elements. Valence patterns underpin technologies that generate power, enable manufacturing, and drive modern society.

Valence Electrons and Chemical Bonds

The octet rule states that atoms bond to obtain eight valence electrons, achieving a stable, noble gas configuration. Atoms can complete an octet by transferring, sharing, or accepting electron pairs during bond formation.

Ionic bonds involve electron transfer to obtain octets. Covalent bonds involve electron sharing to obtain filled valence levels. The specific bonding behavior of an element depends directly on its number of valence electrons.

Valence Bond Theory

Valence bond theory describes covalent bonding based on atomic valence requirements. It states that overlap and mixing of valence atomic orbitals leads to the formation of molecular orbitals. This theory successfully explains variations in bonding patterns.

For example, carbon forms four bonds as each carbon atom overlaps with four other p orbitals. Nitrogen forms three bonds from overlap of its three p orbitals. Oxygen forms two bonds from overlap of its two p orbitals. Observed bonding agrees with predicted orbital overlaps using valence bond theory.

Molecular Orbital Theory

Molecular orbital theory describes bonding in terms of the combination and rearrangement of valence electrons into molecular orbitals. It provides a more detailed quantum mechanical treatment of covalent bonding.

This theory recognizes that atomic orbitals merge into new molecular orbitals encompassing the entire molecule. The number of electrons in these molecular orbitals again depends directly on the number of valence electrons of the constituent atoms.

Both of these major bonding theories emphasize the essential role of valence electrons in dictating bonding behavior.

Electronegativity and Bond Polarity

Differences in electronegativity arise from variations in valence electron configurations. Nonmetals have higher electronegativity values than metals. Noble gases have very low electronegativities.

Bond polarity describes how evenly electrons are shared between atoms. Polar bonds have unequal electron sharing driven by differences in electronegativity. Again, these concepts are intimately tied to valence electron distributions.

Understanding periodic trends in electronegativity and polar bonding relies fundamentally on the concept of valence and its role in chemical bonding.


In summary, the valence of an element can be determined from its electron configuration, periodic table position, reactivity patterns, and bonding behavior. For magnesium, key points include:

  • Magnesium has a valence of 2
  • Its electron configuration is [Ne]3s2 with 2 valence electrons
  • It forms ions like Mg2+ by losing 2 electrons
  • Knowledge of valence allows prediction of chemical formulas

More generally:

  • Metals tend to lose valence electrons, nonmetals gain them
  • Main group elements often have valences matching their group number
  • Transition metal valences depend on d electron configurations
  • Valence determines periodic trends in reactivity

Understanding elemental valence is critical across scientific disciplines. Valence concepts form the basis for theories of chemical bonding. Creative control over valence electrons drives innovations in technology and materials.

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